0.18 Reaction equilibrium in the gas phase  (Page 2/6)

It is important to note that we can vary the relative amount ofNH ₃ produced by varying the temperature of the reaction, the volume of the vessel in which the reaction occurs, or the relative startingamounts of N ₂ and H ₂ . We shall study and analyze this observation in detail in latersections. For now, though, we demonstrate that the concept of reaction equilibrium is general to all reactions.

Consider the reaction

If we begin with 1.00 mole of H ₂ and 1.00 mole of i ₂ at 500 K in a reaction vessel of fixed volume, we observe that, at equilibrium, $n_{HI}=1.72\mathrm{mol}$ , leaving in the equilibrium mixture $n_{H_2}=0.14\mathrm{mol}$ and $n_{I_2}=0.14\mathrm{mol}$ .

Similarly, consider the decomposition reaction

At 298K in a 100.0L reaction flask, 1.00 mol ofN ₂ O ₄ partially decomposes to produce, at equilibrium, $n_{N{O}_2}=0.64\mathrm{mol}$ and $n_{N_2{O}_4}=0.68\mathrm{mol}$ .

Some chemical reactions achieve an equilibrium that appears to be very nearly complete reaction. Forexample,

If we begin with 1.00 mole of H ₂ and 1.00 mole of Cl ₂ at 298K in a reaction vessel of fixed volume, we observe that, at equilibrium, $n_{H\mathrm{Cl}}$ is almost exactly 2.00 mol, leaving virtually no H ₂ or Cl ₂ This does not mean that the reaction has not come to equilibrium. It means instead that, at equilibrium, there are essentially noreactants remaining.

In each of these cases, the amounts of reactants and products present at equilibrium vary as theconditions are varied but are completely reproducible for fixed conditions. Before making further observations that will lead to aquantitative description of the reaction equilibrium, we consider a qualitative description of equilibrium.

We begin with a dynamic equilibrium description. We know from our studies of phase transitions thatequilibrium occurs when the rate of the forward process ( e.g. evaporation) is matched by the rate of reverse process ( e.g. condensation). Since we have now observed that gas reactions also come to equilibrium, we postulate that at equilibrium the forwardreaction rate is equal to the reverse reaction rate. For example, in the reaction here , the rate of decomposition ofN ₂ O ₄ molecules at equilibrium must be exactly matched by the rate ofrecombination (or dimerization ) of NO ₂ molecules.

To show that the forward and reverse reactions continue to happen at equilibrium, we start with theNO ₂ and N ₂ O ₄ mixture at equilibrium and we vary the volume of the flask containing the mixture. We observe that, if we increase the volumeand the reaction is allowed to come to equilibrium, the amount of NO ₂ at equilibrium is larger at the expense of a smaller amount of N ₂ O ₄ . We can certainly conclude that the amounts of the gases atequilibrium depend on the reaction conditions. However, if the forward and reverse reactions stop once the equilibrium amounts ofmaterial are achieved, the molecules would not "know" that the volume of the container had increased. Since the reactionequilibrium can and does respond to a change in volume, it must be that the change in volume affects the rates of both the forward andreverse processes. This means that both reactions must be occurring at equilibrium, and that their rates must exactly match atequilibrium.