Re: Keeping Soda Carbonated

Henry's Law is not going to help with this question.  The concepts of
Chemical Equilibrium are the most useful in this situation.  Carbon
dioxide (CO2) is soluble in H2O.  About 90mL of CO2(gas=g) at 1atm will
dissolve in 100mL of H2O at 20C and about 1% reacts to form carbonic
acid (H2CO3) (ref 1)
			

			~1%
      CO2(g)  +   H2O(l) -------->             H2CO3
		         <-------------------
			~99%

The free gas (CO2) and water are normally favored over the acid.
However, this is an equilibrium situation, not a chemical reaction that
goes to completion.  Modification of the conditions of the system can
shift the equilibrium.  For the case we're interested in, the volume of
the system is fixed (such as a soda bottle) and an increase in the
overall pressure (all gasses) would shift the equilibrium in favor of
the acid. This is true as the release of more CO2 would further increase
the pressure, and the equilibrium only allows for the overall pressure
to go to some finite value, otherwise our soda bottles would
spontaneously explode!!  Or at least get uncomfortably pressurized.
We can artificially push the equilibrium in favor of the acid by
intentionally increasing the pressure of the system (soda bottle) with
an air pump.  By encouraging the formation of the carbonic acid we keep
our previously opened drink from going flat.  When the drink is opened
again the acid we've kept in the liquid decomposes to H2O & CO2 and our
drink remains pleasantly 'fizzy'.  The operating principle is called Le
Chatelier's Principle:  When a disturbance is imposed on a system in
equilibrium, the equilibrium shifts in such a way as to undo, in part,
the effect of the disturbance.

References:  1) B.H. Mahan, University Chemistry, 2nd ed, Adison Wesley,
1969.

2) F. Brescia, J. Arents, H. Meislich, A. Turk, Fundamentals of
Chemistry A Modern Introduction, Academic Press, 1970.

3) J.V. Quagliano, L.M. Vallarino, Chemisrty, 3rd ed.,   Prentice Hall,
1969.