Changing Position of a Gaseous Equilibrium

Description

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• Suppose the reaction below is at equilibrium:

  2 NO₂ N₂O₄ (brown) ΔH = (-57.2 kJ/ mol N₂O₄)

• In this case, the mass action expression for the reaction is written as:

  [N₂O₄ ]e / ([NO₂]e)²

• If this system is at equilibrium, this ratio does not change. Also, changing any conditions of concentration will cause the concentrations to change so that the new equilibrium reached will, when the concentration values are plugged into that expression, give the same numerical value. We call this the equilibrium constant. Since gases are involved in the equilibrium, the concentrations values are volume related.
• Use Le Chatelier's principle to predict the shift in equlibrium. Predict the shift in color with increased temperature. Predict the initial shift in color as volume decreases. Predict the shift in color with time as equilibrium is established at the new smaller volume.

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1. Use a support stand, ring, gauze, Bunsen burner, 150-mL beaker, and 100 mL of water to set up a hot water bath. Prepare an ice water bath using a 150-mL beaker, 50 mL of water, and some crushed ice.
2. Work in a hood. Place 10 mL of concentrated nitric acid in a 50-mL beaker. Place a piece of copper metal in the liquid. Note any evidence of reaction. Note the color of the gas bubbles in the liquid. Note the color of the gas that emerges from the bubbles as they burst in air.
3. Use a plastic or glass syringe with a volume of 5-mL to 30-mL. The gas will corrode the plunger of a plastic syringe; do not plan on reusing. Withdraw a gas sample that half-fills the syringe. Close the tip of the syringe with a plastic cap. Note the intensity of the red-brown color in the syringe.
4. Rapidly compress the gas to about half of the original volume. Note the color intensity change that takes place immediately. Note the change with time shortly thereafter. Repeat the observations of compressing and expanding the volume by moving the plunger quickly and noting the immediate effect and subsequent effect (seconds later) on the color of the gas.
5. Immerse the syringe in the ice water bath. Hold the plunger position fixed. Note the color of the gas immediately after immersing and note any color changes with time.
6. Remove the syringe from the ice water bath and immerse in a hot water bath. Hold the plunger position fixed. Note the color immediately after immersion and note any color changes with time.
7. Remove the syringe from the hot water bath, reimmerse it in the ice bath, and note any changes.

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Name ___________________________ Class ________

Teacher__________________________

DoChem 111 Changing Position of a Gaseous Equilibrium

Watch the movie.

1. Record the color changes observed. Note whether reaction is shifted toward products or reactants.

   2 NO₂ N₂O₄ (brown) ΔH = (-57.2 kJ/ mol N₂O₄)

   Change	Darker or lighter	Products or reactants
   Compress
   Expand
   Heat
   Cool

2. Predict the effect on the color if a volume of air equal to the volume of nitrogen (IV) oxide were added to the syringe and then the system was compressed to the original volume. That is, what happens if an equal volume of air is injected.
3. Design a way to perform the experiment just suggested.

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Purpose

To illustrate the effect of changing temperature and pressure on the position of an equilibrium.

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• fume hood
• tripod
• wire gauze
• Bunsen burner
• matches
• 2 100-mL beaker
• 50-mL beaker
• ice water
• hot water
• copper penny (pre-1982)
• syringe with cap
• nitric acid

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Use a piece of copper foil or wire if no pennies are available.

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Teacher preparation: 30 minutes

Presentation: 20 minutes

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Concentrated nitric acid is corrosive to tissue. Nitrogen (IV) oxide is extremely toxic and corrosive.

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• Do NOT perform this experiment without a hood or an excellent ventilating system that vents out-of-doors.
• Use a hood to avoid inhaling nitrogen (IV) oxide vapors.
• Handle concentrated nitric acid with extreme caution. Wear gloves and apron in addition to safety goggles. Wash spills immediately. Have a working safety shower nearby.

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• Open the syringe under a hood. Rinse the syringe with water several times. Remove the plunger from the syringe, and allow to dry.
• If the plunger of a plastic syringe is not corroded, store for reuse. Cut corroded plastic syringes in half with a sharp knife and discard with ordinary trash.
• Work under a hood. Wear gloves. Once the copper metal has dissolved and the reaction solution has cooled, pour the contents into a 250-mL beaker containing 200 mL of tap water.
• Neutralize excess nitric acid before disposing of the solution at the sink.

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Presentation Question:

• Predict the effect on the color if a volume of air equal to the volume of nitrogen (IV) oxide were added to the syringe and then the system was compressed to the original volume. That is, what happens if an equal volume of air is injected.
  • This change would have no effect on the concentration (partial pressure) of either NO₂ or N₂O₄. Therefore, the prediction is that the equilibrium position will not change.
• Design a way to perform the experiment just suggested.
  • Cover the end of the syringe with a soft plastic cover (such as the one being used). Fill a second syringe with an equal volume of air. Put a sharp needle on the tip of this syringe. Insert the sharp needle of the second syringe through the cover of the first syringe. Hold the first syringe fixed, and press all of the air out of the second syringe.

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• Equilibrium effects are among the most difficult to show students in simple, easily understood systems. It happens that the NO₂ - N₂O₄ equilibrium has just the right properties to illustrate two effects of an equilibrium.
• Suppose the reaction below is at equilibrium:

  2 NO₂ N₂O₄ ΔH = (-57.2 kJ/ mol N₂O₄)

• In this case, the mass action expression for the reaction is written as:

  [N₂O₄ ]e / ([NO₂]e)²

• If this system is at equilibrium, this ratio does not change. Also, we tell students that changing any conditions of concentration will cause the concentrations to change so that the new equilibrium reached will, when the concentration values are plugged into that expression, give the same numerical value. We call this the equilibrium constant.
• Suppose the gaseous system suddenly has its volume halved, as it would in a syringe that is rapidly compressed. Then the new value, before any net reaction takes place, is:

  (2 x [N₂O₄ ]e) / (2 x [NO₂]e)² =(0.5 x [N₂O₄ ]e) / ([NO₂]e)²

• If you compress the system, the new ratio is too small and needs to increase. It will increase if there is a net reaction where NO₂ reacts to form N₂O₄. Compressing the mixture initially causes the color to intensify because the concentration of the colored NO₂ is doubled. However, this scheme predicts that the reaction should show a net formation of colorless N₂O₄ at the expense of NO₂, and the color should fade. This fading is observed. Expansion leads to the opposite effect; the colors immediately fade as the result of dilution, but then increase due to net formation of colored NO₂. This system is effective for a demonstration only because the reaction takes a few seconds. If it took, say 20 seconds, our eyes would not perceive such a gradual change, and a control would be needed. If the reaction were much faster, the equilibrium would be maintained during the compression, and no change would be noted after the compression (or expansion) was complete.
• This reaction is also temperature sensitive and shows an easily interpreted temperature effect. One would guess that a bond formation would release energy. Thus adding energy, perhaps by heating, would shift away from product to colored reactant. Thus the system at constant volume ought to be more colored at high temperature than at low temperature. This is readily observed.

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• Le Chatelier's Principle
• pressure equilibrium
• temperature equilibrium
• equilibrium

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