A simple reaction to demonstrate reversibility is that of the carbonation of water. Water condenses in the atmosphere where CO2(g) exists. Some of this CO2(g) dissolves in the water. Once in solution the CO2(g) reacts with water, H2O, to produce carbonic acid, H2CO3(aq). The amount of CO2(g) dissolving depends directly on the pressure of CO2(g) in the atmosphere. The chemical reaction for CO2(g) and water is shown as (1) below. Bottling companies take advantage of this chemistry to produce carbonated drinks.
Of interest to the students is the fact a given reaction, the proportion found by this methods is always constant when equilibrium is reached. That is to say, no matter where one starts with concentration and/or pressures the final relationship of the product of the amounts of the products divided by the product of the amounts of the reactants will always, always be the same.
Equilibrium is defined as the point where the forward and reverse reactions occur at the same rate. To an observer, the reaction appears to have halted since the observer is viewing at a macroscopic level. At the molecular level, each reaction which makes more products is countered by a reverse reaction (breakdown) of some products into the original reactants. Because there is a balance (equilibrium), the concentration or pressures are no longer changing and a static or constant relationship occurs.
The K for reaction (1) is:
One might note that since [H2O(l)]
is constant it can
be moved to the lefthand side of the equation and included in the
constant for the
CO2(g)/H2CO3(aq) equilibrium
like so:
Note that the magnitude of K is small, 0.0347, which indicates
the reaction equilibrium lies (or favors) the CO2(g)
side of the reaction. One needs a high pressure of
CO2(g) to produce much
H2CO3(aq).
In a bottled drink, the CO2(g) pressure was
artificially high when the bottle was sealed. Thus the reaction
equilibrium is forced to favor the products. When the consumer
opens the bottle, CO2(g) is released from solution.
This loss of CO2(g) is an example of what is known as
Le Chatelier's Principle .
Le Chatelier's Principle
states that a reaction at equilibrium will shift to reattain
equilibirum when a change is made to one of the components of the
reaction. In the case of the bottle, the
PCO2(g) is high in the small volume of
space above the drink. This pressure change is noted by the
"pffft" that occurs when a bottle (or can) of carbonated drink is
opened. Once the drink is open the PCO2(g)
drops to atmospheric levels and the CO2 dissolved in
the water begins to flow back out. The
H2CO3(aq) begins to form CO2(g)
and H2O(l). Reaction (1) is running in reverse.
There is a rate effect which is why the soda doesn't go "flat "
immediately but the equilibrium constant above shows that ratio
of the concentration of H2CO3(aq) to
pressure or CO2(g) will be 0.0347 M/atm when
equilibrium is reestablished. M is the symbol for
molarity.
[H2CO3(aq)]
The [H2O(l)],
color used for clarity only and square brackets, [ ], are used to
symbolize concentration, is constant and is not used in
calculating a K for this reaction. If water were in limited
amounts, one would include its concentration in the calculation.
The K for that reaction would be different than the one given
here.
K = -----------
PCO2[H2O(l)]
[H2CO3(aq)]
And we can replace the K.[H2O(l)]
with another K where K represents the new constant. I am using colors only to
indicate that the Ks are different.
K.[H2O(l)] = ---------
PCO2
[H2CO3(aq)]
For reaction (1), the above constant has a value of 3.47 x
10-2. The units of this constant will be molarity
over atmosphere. Some K's have no units while others have mixed
units. This is wholly dependent upon the reaction. It is the
magnitude (size) of the K that is important, not the units.
K = -----------
PCO2