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					      Lab 2 Disturbing an Equilibrium
      Introduction
      In this lab activity, you will use Le Châtelier's Principle to predict the effect of imposed
      changes on systems at equilibrium. Then you will test your predictions by disturbing a
      system at equilibrium, observing the effect of the disturbance, and interpreting the
      observation in terms of Le Châtelier's Principle.


      Problem
      Do systems at equilibrium respond to imposed changes in predictable ways?


      Pre-Lab Predictions
      Create a table like this one. Complete the third column of Table 1 by writing "shifts left",
      "shifts right" or "no change" to indicate the way you think the position of the equilibrium
      will be affected by the imposed change. You will record your observations in the fourth
      column as you perform steps in the procedure below.

      Table 1: Predicting changes in the positions of equilibria.
                                                 Imposed       Pre-Lab
System                                                                         Observations
                                                 Change        Predictions
                                                   Increase
                                                     [H+]
                                                   Increase
Chromate-Dichromate
                                                     [OH-]
                                                   Decrease
                                                    [OH-]
                                                   Decrease


Phenolphthalein ion complex                        Increase
                                                     [H+]
                                                   Decrease
                                                     [H+]
Cobalt(II) ion Complex                             Add Heat

                                                   Decrease
                                                     Heat
                                                   Add Heat
NO2-N2O4 Gas Phase
                                                   Decrease
                                                     Heat
                                                   Increase
                                                   Pressure
Materials
Part 1
   • 5 test tubes
   • a 24 well plate
   • four clean medicine droppers
   • 0.10 M potassium chromate
   • 1.0 M hydrochloric acid
   • 1.0 M sodium hydroxide
   • 1.0 M iron(III) nitrate (ferric nitrate)
   • 1.0 M barium chloride

Part 2
   • 50 mL beaker
   • distilled water
   • graduated cylinder
   • 1.0 M sodium hydroxide
   • phenolphthalein indicator solution
   • 1.0 M hydrochloric acid

Part 3
   • 50 mL beaker
   • graduated cylinder
   • ethanol
   • metal scoopula
   • cobalt(II) chloride
   • stirring rod
   • hydrochloric acid (concentrated if available)
   • 2 test tubes
   • ice water bath (large beaker of ice water)
   • hot water bath (large beaker of very hot kettle water

Part 4
   • no materials - refer to photographs on pages 523-524 of MHR textbook


Procedures
Safety Precautions
You will be using several dangerous and hazardous substances in this activity. Wear
protective eyewear and clothing throughout the activity. Specific hazards include
corrosive hydrochloric acid, caustic sodium hydroxide, poisonous chromate and barium
compounds, and flammable ethanol. If your skin comes in contact with any of the
chemicals during this activity, rinse the affected area thoroughly under cold running
water for at least 15 minutes and notify your supervising teacher of the incident.
Part 1 - Effects of Concentration Changes
1. Transfer small amounts hydrochloric acid, sodium hydroxide, iron(III) nitrate, and
   barium chloride into wells A1, B1, C1, and D1 of the 24 well plate.
2. Label five test tubes as A, B, C, D, and E. Pour 5 mL of potassium chromate solution
   into each of five test tubes. Set tube E aside as your control.
3. Using a clean medicine dropper, add five drops of HCl to test tube A. Compare the
   colour of the solution in test tube A to the control. Record your observations.
4. Using a clean medicine dropper, add five drops of HCl to test tube B, then using a
   second clean medicine dropper, add five drops of NaOH to test tube B. Compare the
   colour of the solution in tube B to the control before and after the addition of NaOH.
   Record your observations.
5. Using a clean medicine dropper, add 5 drops of Fe(NO3)3 to test tube C. A
   precipitate, Fe(OH)3, should form. Set the tube aside for a few minutes to let the
   precipitate settle. Compare the colour of the solution above the precipitate to the
   colour of the control. Record your observations.
6. Using a clean medicine dropper, add five drops of BaCl2 to test tube D. A precipitate,
   BaCrO4, should form. Set the tube aside for a few minutes to let the precipitate
   settle. Compare the colour of the solution above the precipitate to the colour of the
   control. Record your observations.
7. Reserve the well plate contents for part 2.
8. Pour the contents of all tubes into a sink and flush with plenty of water. Clean all
   glassware and plastic ware thoroughly.



Part 2 - Effects of Concentration Changes II
1. Top up HCl and NaOH solutions in well plate if necessary.
2. Pour about 19 mL of distilled water into the 50 mL beaker and add 1 mL of 1.0 M
   sodium hydroxide solution.
3. Add one or two drops of phenolphthalein indicator solution and swirl the beaker
   gently. The colour you observe is caused by a complex ion represented by the short-
   hand symbol Ph-.
4. Using a clean medicine dropper, add HCl (hydrogen ions) to the beaker one drop at
   a time until a colour change occurs. Record/describe the colour change in your data
   table. The HPh form of the indicator is colourless.
5. Using a clean medicine dropper add NaOH to the beaker until you detect another
   colour change. Record/describe the colour change in your data table.
6. Pour the contents of the beaker and the well plate into a sink and flush with plenty of
   water. Clean all glassware and plastic ware thoroughly.
Part 3 - Effect of Temperature I
1. Measure 15 mL of ethanol and transfer to a small (50 mL) beaker.
2. Use a metal scoopula to transfer a pea-sized amount of CoCl2 into the beaker
   containing ethanol. Stir the CoCl2 to dissolve using a stirring rod.
3. If the solution is pink, add drops of concentrated HCl (if available) or saturated
   sodium chloride solution (if concentrated HCl is not available) to turn it to a blue or
   purplish colour.
4. Divide the solution among the two test tubes.
5. Place one tube in a hot water bath and the other in a ice water bath. Record the
   colour of each solution.
6. Switch the tubes around so that each is in a different bath. Record the colour of each
   solution.
7. Pour the contents of the tube into a sink and flush with plenty of water. Clean all
   glassware and plastic ware thoroughly.

Part 4 - Effect of Temperature and Effect of Pressure
1. Record the colour of the gas mixture at 0°C (left picture p.523 MHR).
2. Record the colour of the gas mixture at 100°C (right picture p.523 MHR).
3. Record the colour of the mixture at normal atmospheric pressure (left picture, p.524
   MHR)
4. Compare the colour of the mixture at 2X atmospheric pressure before the system
   responds to the change (middle picture p.524) and after the systems response to
   increased pressure (right picture p.524 MHR).


Analysis and Conclusion

Effect of Changes in Concentration

   1.    List the tests which involved increasing the concentration of a reactant. Did
         your observations indicate a shift to the left or to the right? In terms of
         concentration changes, account for the observed shifts in positions of the
         equilibria.

   2.    List the tests which involved increasing the concentration of a product. Did
         your observations indicate a shift to the left or to the right? In terms of
         concentration changes, account for the observed shifts in positions of the
         equilibria.

   3.    List the tests which involved decreasing the concentration of a reactant. Did
         your observations indicate a shift to the left or to the right? In terms of
      concentration changes, account for the observed shifts in positions of the
      equilibria.

4.    List the tests which involved decreasing the concentration of a product. Did
      your observations indicate a shift to the left or to the right? In terms of
      concentration changes, account for the observed shifts in positions of the
      equilibria.

5.    State a generalization about increasing the concentration of a species in a
      chemical equilibrium.

6.    State a generalization about decreasing the concentration of a species in a
      chemical equilibrium.

Effect of Changing Temperature

Consider these equations:

                               N2O4    (g)   +   59 kJ   2 NO2   (g)

       Co(H2O)2 +
              6 (aq)    +      4     −
                                   Cl(aq)    +   50 kJ   CoCl2 −
                                                             4(aq)     +   6 H2O (l)

7.    Are the forward reactions endothermic or exothermic? Why?

8.    How did these equilibria respond to an increase in temperature?

9.    Provide an explanation for each observed shift in terms of rates of reaction and
      activation energy.

10.   State a generalization about the effect of increasing the temperature of a
      system on the position of an equilibrium.

Effect of Pressure

11.   What effect did reduction of the volume of the syringe (as pictured on page 524
      of MHR) have on the pressure of the system?

12.   Based on the illustrations on page 524 of MHR, how did the equilibrium
      position respond to the change in pressure?

13.   How did the amounts of the two gases change as a result of the pressure
      change?

14.   Predict the effect of a reduction in the volume of this system:
      2 IBr(g)         I2(g)       +   Br2(g)
        on the equilibrium position and the mole amount of each gas species.

  15.   State a generalization about the effect of a pressure change due to a change
        in volume on the position of an equilibrium.



Conclusion

  1.    Which of the predictions you made using Le Chatelier’s principle differ from the
        observations you made in each part of the activity?

  2.    What are the possible causes of the differences?

				
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