Aim: To illustrate
how, as a result of the equilibrium law, the equilibrium mixture is changed, after
similarly influencing the concentrations of all components.
Equilibria in the gas phase
can be shifted by changing the volume of the closed system. The concentrations of all the
components are thereby changed in the same way. Illustration E10a shows the evolution of
the concentrations of NO2 and N2O4 as a function of time,
after decreasing the volume of the closed system. As with the external addition of NO2,
the starting point is the equilibrium mixture resulting from Experiment A. At time t1
the reaction volume is halved, for example by pushing in a piston in a closed cylinder.
This doubles the concentrations of NO2 and N2O4
(resulting in concentrations of 0.022 mol/L and 0.054 mol/L respectively), which no longer
comply with the equilibrium conditions at 25 °C.
This abrupt doubling in N2O4- and NO2-concentrations is
counteracted by concentration changes in the reaction mixture, the N2O4-concentration
increasing slowly to 0.057 mol/L and the NO2-concentration slowly decreasing
from 0.022 mol/L to 0.016 mol/L. This is, of course, consistent with the 2:1 stoichiometry
of the reaction.
The fact that the N2O4-concentration further increases, poses the
question: how is this to be explained? Has the reaction stoichiometry changed or is it
simply because the N2O4-concentration was the greater? The actual
explanation is contained in the equilibrium law.
The green band in
illustration E10b shows the variation in Qc as a function of time. The
doubling of the concentrations reduces Qc from 222 to 111, which is
half the Kc-value of 222. The equilibrium law is relentless, the
disturbance has to be counteracted by changing the concentrations rapidly or slowly to the
new values dictated by the Kc-value. These changes must result in an
increase in the numerator and hence a corresponding decrease in the denominator, as
dictated by the stoichiometry of the reaction. The new equilibrium concentration of N2O4
is, at 0.057 mol/L slightly higher than the concentration resulting from the volume
decrease. The new equilibrium concentration of NO2 is, at 0.016 mol/L slightly
lower than the concentration resulting from the volume decrease.
A calculation and a table such as in
illustration E09b can be useful in understanding the relationship between the change in Qc
(the green band) and the changes in
N2O4 (white band)-and NO2 (brown band)-concentrations as
a function of time.
If the equilibrium concentrations after the volume decrease, are compared with those
before the volume decrease, the concentrations of both components have increased. This is
solely due to the volume decrease. In the case of N2O4, with the
lower stoichiometric number, the change in concentration becomes even greater, whereas in
the case of NO2 the change has been counteracted as much as possible. This is
again in accordance with the general principle of the smallest change, the final
concentration of NO2 not being able to dip below the previous equilibrium
In the case of simultaneous and uniform concentration change of reactants and products,
the equilibrium law can again be relied upon to predict the composition of the new
equilibrium mixture. It is a simple rule of thumb that if the concentrations of the
components are simultaneously and uniformly increased, then according to the
Chateliers principle, the equilibrium will shift to the side of the reaction
equation with the fewer molecules to keep the number of molecules per unit volume in the
system as constant as possible. In the event of simultaneous and uniform concentration
decrease, on the other hand, the equilibrium will shift to the side of the reaction
equation with more molecules, again to keep the number of molecules per unit volume in the
system as constant as possible. In general in the event of volume change, its effects must
be predicted in terms of concentration changes.