Equilibrium Shift on Heating: Interpreting Colour and Yield Data
A sealed flask contains the equilibrium mixture 2NO2(g) ⇌ N2O4(g), which is brown (NO2) and colourless (N2O4). The forward reaction is exothermic, ΔH = −57 kJ/mol. A student heats the flask from 25 °C to 60 °C and observes the brown colour deepen. A second experiment adds more N2O4 at constant volume and temperature; the mixture darkens briefly, then lightens. A third experiment compresses the mixture; the colour first deepens, then lightens. The student concludes that "heating always favours the products" and that "adding a reactant always increases product yield."
1. Use Le Chatelier's principle and the equilibrium constant to explain the effect of heating.
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Le Chatelier reasoning. The forward reaction is exothermic (ΔH = −57 kJ/mol), so heat can be regarded as a product. Raising the temperature adds "product," so the system shifts in the endothermic direction - the reverse reaction - to absorb heat. The equilibrium shifts left, increasing [NO2] and decreasing [N2O4]. The deeper brown colour confirms more NO2, so the student's claim that heating favours products is wrong for this reaction.
Equilibrium constant reasoning. For an exothermic reaction, K decreases as temperature rises. Since K = [N2O4]/[NO2]², a smaller K means the ratio of product to reactant falls, i.e. the position shifts toward NO2. This is the more rigorous statement: temperature changes the value of K itself, whereas concentration and pressure changes only move the position along a fixed K.
Conclusion. Heating shifts this equilibrium toward reactants, consistent with the observed darkening. The student's generalisation fails because the direction of the temperature effect depends on the sign of ΔH, not on a universal rule.
2. Explain the two-step colour changes when (a) N2O4 is added and (b) the mixture is compressed.
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(a) Adding N2O4 at constant volume and temperature. Immediately after addition, [N2O4] rises, so the mixture should become lighter (more colourless molecules, though some NO2 is still present). The brief darkening is a transient: the sudden increase in [N2O4] temporarily increases the rate of the reverse reaction (N2O4 → 2NO2), producing extra NO2 before the system re-equilibrates. As equilibrium is re-established, some of the added N2O4 is converted to NO2, so [NO2] ends up higher than before but lower than the transient peak; the net colour is intermediate. The key point is that adding a reactant does not simply increase product yield - it shifts the position, but the final [NO2] depends on the new equilibrium, and K is unchanged at constant temperature.
(b) Compression. Reducing the volume raises the total pressure. The system shifts toward the side with fewer gas moles - the product side (2 mol gas → 1 mol gas) - so the equilibrium shifts right, decreasing [NO2] and lightening the colour. The initial deepening is simply the concentration effect of squeezing the same moles into a smaller volume before the equilibrium adjusts. The student's claim that adding a reactant always increases yield is too simplistic: the outcome depends on the equilibrium position and on whether K changes.