Materials Included In Kit

Additional Materials Required

Prelab Preparation

Sealed Pipets
The following steps should only be performed in an efficient, well-operating fume hood. Wear chemical splash goggles and chemical-resistant gloves and apron. Place about 20 mL of concentrated nitric acid in a large (500-mL) Erlenmeyer flask. Cut a small piece of copper foil (about 1" x 1") and add the piece to the nitric acid. Note: Two 0.5" x 0.5" pieces of copper may be used in place of a 1" x 1" piece. Wait a few minutes while the acid oxidizes the copper and brown nitrogen dioxide fumes appear in the flask. Squeeze as much of the air as possible out of a jumbo polyethylene pipet bulb, place the pipet in the neck of the flask, and release the squeeze to draw the nitrogen dioxide gas into the pipet. Heat the stem (long end) of the pipet in a Bunsen burner flame and seal the stem shut with a pair of pliers. Note: The stem of the pipet will turn clear when heated. Do not let it catch fire. Remove the pipet from the flame and seal it. The 1" x 1" piece of copper will produce enough brown gas to fill about 16 pipets. Once all the copper has reacted, add another 1" x 1" piece of copper to the flask. Repeat this process until all the pipet bulbs are filled.

Safety Precautions

Nitrogen dioxide is a highly toxic gas. The gas will be supplied to the students in sealed polyethylene pipet bulbs. Instruct the students not to cut the pipet bulbs or puncture them in any way. Students should not continue with the procedure if any breaks or tears in the bulbs are seen or yellow stains appear on students’ hands. Instruct the students to immediately take any leaking pipets to the fume hood and notify you. Have them wash hands thoroughly with soap and water. The polyethylene pipets may be slightly permeable with respect to nitrogen dioxide and the gas may slowly leak out of the pipets. The resulting pressure decrease may cause the pipets to pucker over time. Prepare fresh sealed pipets as needed. Wear chemical splash goggles and chemical-resistant gloves and apron. Wash hands thoroughly with soap and water before leaving the laboratory. Please consult current Safety Data Sheets for additional safety, handling and disposal information.

Disposal

Consult your current Flinn Scientific Catalog/Reference Manual for general guidelines and specific procedures, and review all federal, state and local regulations that may apply, before proceeding. The gas-filled pipets should be placed in an efficiently operating hood immediately after use. To dispose of the gas, cut the ends of the sealed pipet bulbs and immerse the bulbs in water in a large, 2-L beaker. Allow the gas to dissolve in and react with the water overnight. The resulting acidic aqueous solution may be neutralized and disposed of down the drain with plenty of excess water according to Flinn Suggested Disposal Method #24b.

Lab Hints

• Treat this microscale lab as a mini-lab alternative to lecture. If the hot and cold water baths are prepared in advance, the laboratory work can reasonably be completed in 20–30 minutes. This allows ample time within a typical classroom period to review the background material, answer the Prelaboratory questions, make predictions, and discuss the results. This microscale experiment may also be performed as a demonstration using a document camera connected to a TV monitor. Alternatively, large-scale, sealed glass diffusion tubes filled with nitrogen oxides are available (Flinn Catalog No. AP8476) for demonstration purposes.
• Enough bulbs are available to give each student group three bulbs to work with. One bulb should be reserved as a control (reference) to compare the observed color changes.
• The color changes observed upon heating and cooling the pipets are very pronounced. The color changes observed when the pipets are squeezed are not as dramatic and may, in fact, be contradictory. Use the effect of pressure as a “teachable” moment to show students that the properties of a gas are interrelated.
• When the pipets are squeezed to reduce the volume, the color at first darkens, then appears to gradually lighten. Pictures of this phenomenon using sealed syringes are included in many textbooks, and the results are usually described in terms of a slow relaxation time to reach equilibrium. According to the textbooks, the color becomes darker as the volume decreases because the concentration of NO₂ increases. Slowly, however, according to these sources, the increase in pressure causes the position of equilibrium to shift in favor of N₂O₄, because there are fewer gas molecules on the product side than on the reactant side. This is a kinetic argument—it takes time to reach equilibrium.
• According to kinetic studies carried out using special techniques, the dimerization of NO₂ is extremely fast even at room temperature. Calculations based on the rate data suggest that the time needed to reach equilibrium is on the order of microseconds—significantly faster than the time needed to squeeze a pipet bulb and observe the color change. It has been argued, therefore, that the immediate darkening observed when the pressure increases is actually due to temperature. Compressing the gas into a smaller volume increases the temperature as well as the pressure. The resulting temperature increase shifts the equilibrium in favor of NO₂. The color then gradually fades as the gas mixture cools to room temperature and the pressure increase shifts the equilibrium in favor of N₂O₄. The slow step is not the time needed to reach chemical equilibrium, but rather the time needed to reach thermal equilibrium with the surroundings. If students are holding the pipets in their hands, heat will also be added from their hands. See the article “Approaching Equilibrium in the N₂O₄–NO₂ System: A Common Mistake in Textbooks” (J. Chem. Ed., 2000, 77, 1652–1655) for a discussion of the effects of pressure, temperature, and concentration on the color of the N₂O₄–NO₂ system.

Teacher Tips

• The balanced chemical equation for the formation of nitrogen dioxide is shown.

  Cu(s) + 4HNO₃(aq) → Cu(NO₃)₂(aq) + 2NO₂(g) + 2H₂O(l)

• As discussed in the Background section, nitrogen oxides are primary pollutants responsible for the production of ozone in the lower atmosphere. Nitrogen oxides are also involved in the decomposition of ozone in the upper atmosphere. Paul Crutzen shared the Nobel Prize in chemistry in 1995 for his work on the role of NO and NO₂ as catalysts in the destruction of the ozone layer. There is a lot of interesting gas-phase chemistry of nitrogen oxides going on in the atmosphere.

Answers to Prelab Questions

1. Draw Lewis electron dot structures for the nitrogen oxides mentioned in the Background section: nitric oxide NO, nitrogen dioxide NO₂, and dinitrogen tetroxide N₂O₄.
   {13959_PreLabAnswers_Figure_1}
   Note: The charges shown are formal charges on the atoms. The molecules themselves do not carry an overall charge. Some textbooks represent bonds between atoms having formal l+ and l– charges using arrows for “coordinate–covalent” bonds.
2. Use the electron dot structures of NO and NO₂ to explain why these molecules are considered highly reactive.

   It is impossible to draw Lewis structures for either NO or NO₂ in which all the electrons are paired. A single unpaired electron must be present on either the nitrogen atom or one of the oxygen atoms. Molecules containing unpaired electrons are generally considered more reactive than molecules containing all paired electrons.

3. Although both N₂ and O₂ are naturally present in the air we breathe, high levels of NO and NO₂ in the atmosphere occur mainly in regions with large automobile or power plant emissions. The equilibrium constant for the reaction of N₂ and O₂ to give NO is very small. The reaction is, however, highly endothermic, with a heat of reaction equal to +180 kJ (Equation 7).
   {13959_PreLabAnswers_Equation_7}
   1. Use Le Chatelier’s Principle to explain why the concentration of NO at equilibrium increases when the reaction takes place at high temperatures.

      According to Le Chatelier’s Principle, increasing the temperature shifts the equilibrium in the direction in which heat is absorbed, that is, in favor of NO formation.

   2. Use Le Chatelier’s Principle to predict whether the concentration of NO at equilibrium should increase when the reaction takes place at high pressures.

      According to Le Chatelier’s Principle, increasing the pressure should not affect the position of equilibrium for the reaction, since there are an equal number of gas molecules on each side of the equation. Note: There is also a kinetic argument that can be made. Reactions of gases generally occur much faster at elevated temperatures and pressures.

Sample Data

Answers to Questions

1. Write the chemical equation for the reaction of NO₂ to form the dimer N₂O₄. Include the color of each compound underneath its formula.
   {13959_Answers_Reaction_1}
2. What color change was observed when the gas was cooled? In what direction did the equilibrium shift?

   The color changed from brown to almost colorless when it was cooled. The equilibrium was shifted in the forward direction, that is, in favor of products.

3. What color change was observed when the gas was heated? In what direction did the equilibrium shift?

   The color changed from brown to very dark brown when it was heated. The equilibrium was shifted in the reverse direction, that is, in favor of reactants.

4. Are both reactant and product gases present in the original equilibrium mixture at room temperature? Explain.

   The reversible color changes upon heating and cooling demonstrate that both reactants and products must be present in the original equilibrium mixture at room temperature. We know there are NO₂ molecules present because of the color. We know there are N₂O₄ molecules present because the color got darker when it was heated—some of the N₂O₄ molecules present at room temperature must have dissociated upon heating.

5. Use the results of the heating and cooling experiments to decide whether the dimerization reaction of NO₂ is endothermic or exothermic. Rewrite the chemical equation for the reaction to include the heat term on the reactant or product side, as needed.

   The fact that the reaction shifted in the reverse direction upon heating means that heat is absorbed in the reverse direction. Therefore, the forward reaction is exothermic and heat should appear on the product side of the chemical equation.

   {13959_Answers_Reaction_2}
   Note: This should make intuitive sense. Bond formation is always an exothermic process. In the forward direction, a new bond is formed between the two nitrogen atoms.
6. Use Le Chatelier’s Principle to explain the effect of temperature on the gas phase equilibrium involving NO₂ and N₂O₄.

   According to Le Chatelier’s Principle, increasing the temperature should shift the equilibrium in favor of the reaction that will absorb some of the excess heat that has been added to the system. The opposite argument may be made for decreasing the temperature. In this case, the reaction should shift in favor of the reaction that will release heat.

7. Write the equilibrium constant expression (mass action expression) for the nitrogen oxide equilibrium. Does the value of the equilibrium constant depend on temperature?
   {13959_Answers_Equation_8}
   The value of the equilibrium constant must depend on temperature, since the relative amounts of reactants and products changed when the temperature was changed, even though no additional materials were added to the system. Note: The equilibrium constant may also be expressed in terms of the partial pressures of the gases.
   {13959_Answers_Equation_9}
   The calculated value of K_p at 293 K is 10.5 (based on DG° = –4.84 kJ/mole).
8. According to Boyle’s Law, what happened to the pressure inside the bulb when the bulb was squeezed to half its original volume? Use Le Chatelier’s Principle to predict how this pressure change should affect the position of equilibrium for the NO₂–N₂O₄ reaction.

   According to Boyle’s Law, volume and pressure are inversely related. The pressure inside the bulb increased when the applied volume was reduced. Increasing the pressure in the NO₂–N₂O₄ mixture should shift the equilibrium in favor of the side containing fewer gas molecules, that is, to the product side (one N₂O₄ molecule is formed by the combination of two NO₂ molecules).

9. Discuss the color changes observed when the gas volume was reduced. Do the changes agree with the prediction made above for the effect of pressure?

   When the bulb was squeezed, the gas mixture at first darkened, suggesting that more NO₂ molecules were being formed. The color then gradually faded, but did not go colorless. This does not agree with the prediction made above based on Le Chatelier’s Principle.

10. What other factors or conditions might have influenced the color changes observed when the bulb was squeezed? Hint: Did any of the other gas variables (P, V, T, n) change?

    Squeezing the bulb reduces the volume and thus increases the concentration of the gas molecules if no other changes occur. Compressing the gas molecules inside the pipet bulb also increases the effective temperature of the gas. The temperature increase and pressure increase have opposite effects on the equilibrium. Note: The final color of the gas is slightly darker than in the original pipet bulb. Equilibrium constant calculations show that after thermal equilibrium has been reached, the concentration of NO₂ molecules will be higher even though their partial pressure has been reduced as a consequence of Le Chatelier’s Principle.

References

This experiment has been adapted from Flinn ChemTopic™ Labs, Volume 15, Equilibrium, Cesa, I., Ed., Flinn Scientific, Batavia, IL, 2003.

Introduction

Many important reactions that take place in the atmosphere involve equilibrium concentrations of gas phase reactants and products. What variables affect the position of equilibrium for reactions in the gas phase?

Concepts

• Chemical equilibrium
• Le Chatelier’s principle
• Gas phase reactions
• Nitrogen oxides

Background

Burning fossil fuels for energy “drives” our society and our economy. It is also a major source of environmental concerns and challenges. The release of large amounts of carbon dioxide from the combustion of oil and gas, for example, is a subject of controversy because of its possible contribution to global warming. In addition to carbon dioxide, burning fossil fuels also produces a variety of sulfur and nitrogen oxides. Sulfur oxides are formed via the oxidation of sulfur-containing impurities in coal and oil (Equations 1 and 2) and are a major cause of acid rain. Nitrogen oxides are formed when nitrogen and oxygen—the main components of air—combine with one another in car engines, power plants, or in car exhaust (Equations 3 and 4). Nitrogen oxides are a major component of photochemical smog and air pollution.

As can be seen from Equations 2–4, most of the gas phase reactions that take place in the atmosphere are reversible reactions. Conditions that affect the position of equilibrium for gas phase reactions are therefore of enormous importance in determining the environmental impact of burning fossil fuels. In this experiment, we will consider the properties of nitrogen dioxide and investigate how the principles of equilibrium apply to its reactions.

Nitrogen dioxide (NO₂) is a toxic, reddish-brown gas with an irritating odor. It is primarily responsible for the brownish haze that hangs over many of the world’s largest cities due to air pollution. Nitrogen dioxide is also quite reactive. In the presence of sunlight, for example, it undergoes a light-induced “photochemical” reaction to produce ozone (Equation 5). High levels of nitrogen oxides in the atmosphere are associated, therefore, with high ozone levels as well. The high reactivity of nitrogen dioxide means that it reacts even with itself—two molecules of NO₂ combine to form the “dimer,” dinitrogen tetroxide, N₂O₄, which is a colorless gas at room temperature. Formation of N₂O₄ is a reversible reaction (Equation 6) and quickly reaches a position of equilibrium. The relative amounts of NO₂ and N₂O₄ present at equilibrium depend on pressure and temperature, according to Le Chatelier’s Principle.

Experiment Overview

The purpose of this experiment is to study the effects of changing the temperature and pressure on the relative amounts of NO₂ and N₂O₄ in a sealed tube at equilibrium. Le Chatelier’s Principle predicts how a change in conditions will affect the equilibrium for a reversible chemical reaction—the reaction will shift in a direction that tends to reduce the effect of the imposed change. The effect of changing the temperature depends on whether the reaction is exothermic or endothermic as written, while the effect of changing the pressure depends on the number of gaseous molecules on the reactant versus product side of the reaction equation.

Materials

Prelab Questions

1. Draw Lewis electron dot structures for the nitrogen oxides mentioned in the Background section: nitric oxide NO, nitrogen dioxide NO₂, and dinitrogen tetroxide N₂O₄.
2. Use the electron dot structures of NO and NO₂ to explain why these molecules are considered highly reactive.
3. Although both N₂ and O₂ are naturally present in the air we breathe, high levels of NO and NO₂ in the atmosphere occur mainly in regions with large automobile or power plant emissions. The equilibrium constant for the reaction of N₂ and O₂ to give NO is very small. The reaction is, however, highly endothermic, with a heat of reaction equal to +180-kJ (Equation 7).
   {13959_PreLabAnswers_Equation_7}
   1. Use Le Chatelier’s Principle to explain why the concentration of NO at equilibrium increases when the reaction takes place at high temperatures.
   2. Use Le Chatelier’s Principle to predict whether the concentration of NO at equilibrium should increase when the reaction takes place at high pressures.

Safety Precautions

Nitrogen dioxide is a highly toxic gas. The gas will be supplied in sealed polyethylene pipet bulbs. Do not cut the pipet bulbs or puncture them in any way. Do not continue with the procedure if you see any breaks or tears in the bulbs or if you see yellow stains on your hands. Wash hands at once. Immediately take any leaking pipets to the fume hood and notify your teacher. Wear chemical splash goggles and chemical-resistant gloves and apron. Wash hands thoroughly with soap and water before leaving the laboratory.

Procedure

1. Fill two beakers (250- or 400-mL) half-full with tap water. Heat one beaker on a hot plate to prepare a hot-water bath (75–80 °C) for use in step 4. Add ice to the second beaker to prepare an ice-water bath (0–5 °C) for use in step 5.
2. Measure the room temperature and record it in the data table.
3. Obtain two sealed pipet bulbs filled with nitrogen oxides from your teacher. Observe and record the color of the gas at room temperature.
4. Using forceps or tongs, place one pipet bulb in the hot-water bath for 2–3 minutes. Measure the temperature of the bath and observe the color of the gas. Record this data in the data table.
5. Use forceps or tongs to remove the pipet bulb from the hot-water bath, then immerse the bulb in the ice-water bath. Measure the temperature of the bath and observe the color of the gas. Record this data in the data table.
6. Alternate immersing the pipet bulb in the hot-water and ice-water baths. Are the color changes repeatable? Record all observations in the data table.
7. Place the bulb on a piece of white background paper. Does the gas return to its original color when the bulb returns to room temperature?
8. Take the second pipet bulb and hold it vertically at one end. Squeeze on the bulb and bend it over to compress the gas into a smaller volume. Try to squeeze the gas into about one-half its original volume.
9. Observe and record any immediate color changes that occur when the gas is compressed.
10. Continue squeezing the pipet bulb in this manner for 2–3 minutes. Observe any further color changes that may occur. Compare the color of the gas against a white background with that in the first pipet bulb, which should be at room temperature.
11. Return all pipet bulbs to your teacher for disposal.

Student Worksheet PDF

13959_Student1.pdf