The Sealed Tube That Would Not Follow the Rule
During a unit on equilibrium, a teacher prepares three sealed glass tubes containing an NO2/N2O4 mixture. The tube in ice water is pale brown; the room-temperature tube is medium brown; the tube in hot water is dark brown. Students explain the temperature effect correctly. The teacher then compresses a gas sample in a 100 mL syringe from 60 mL to 30 mL at room temperature. Students observe an immediate darkening, followed by a slight fading over about 30 seconds. A student says, 'The fading proves the equilibrium shifts left, so compression always makes less NO2.' Another student predicts that injecting helium into a rigid, sealed flask at constant volume will also shift the equilibrium toward N2O4 because 'the pressure goes up.' The teacher must decide how to respond in the moment and what follow-up would test whether students are using partial pressures or just total pressure.
1. Evaluate the helium prediction. Does adding helium at constant volume shift the NO2/N2O4 equilibrium? Explain using partial pressures or Q versus K.
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At constant volume and temperature, adding helium raises the total pressure but does not change the partial pressures of NO2 or N2O4, assuming ideal behavior. The equilibrium expression is Kp = P(N2O4)/[P(NO2)]^2. Helium does not appear in the expression, and the mole fractions of the reacting gases decrease only because helium is added; their partial pressures remain the same because their moles and the volume are unchanged. Therefore Q still equals K and there is no shift. The student is conflating total pressure with the partial pressures that determine Q. A different case is adding helium at constant pressure, which increases volume and lowers the reacting gases' partial pressures, shifting toward the side with more gas moles, NO2. The teacher should make that contrast explicit; otherwise students memorize 'pressure shifts to fewer moles' without specifying what pressure means.
2. Interpret the syringe observation. Why does the color darken first and then fade slightly, and how does this sequence relate to concentration and Le Châtelier's principle?
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Compression from 60 mL to 30 mL halves the volume, so the concentrations of both NO2 and N2O4 initially double. The immediate darkening is mainly due to the higher NO2 concentration, not an equilibrium shift. After that, the system responds to the increased partial pressures by shifting toward the side with fewer gas moles, N2O4: 2 NO2 ⇌ N2O4. This consumes some NO2, so the color fades slightly from its darkest instant. The fading does not mean compression makes less NO2 than before compression; the new equilibrium NO2 concentration is still higher than the original because the initial increase was only partially offset. Le Châtelier's principle predicts the direction of the partial offset, not a reversal of the disturbance. This is a useful distinction between concentration effects and equilibrium shifts.
3. Design a short formative assessment or demonstration that would reveal whether students are confusing total pressure with the partial pressure of the reacting gases.
Read the worked answer
Use a paired prediction task. Show two rigid flasks of equal volume at the same temperature, each with the same NO2/N2O4 mixture. Flask A has only the equilibrium mixture; flask B has the same mixture plus helium to double the total pressure. Ask students to predict whether the color of B will be lighter, darker, or the same as A, and to justify with a particle diagram and a Q/K comparison. Include a second pair in which helium is added at constant pressure so the volume doubles. The contrast separates constant-volume from constant-pressure reasoning. A good response says B is the same at constant volume because partial pressures of NO2 and N2O4 are unchanged, while the constant-pressure case shifts toward NO2 because partial pressures fall. Follow with a quick calculation of Q using partial pressures to confirm whether students can connect the symbolic expression to the demonstration.