The instability of hydrated carbon dioxide or "carbonic acid"
The situation is described by the following equilibrium
$$\ce{H2O + CO2 &lt;=&gt; HO-CO-OH}$$
In a closed system, using Le Chatelier’s principle you can manipulate the equilibrium in either direction (e.g. pump in carbon dioxide and increase the concentration of carbonic acid, etc.).  However in an open system, as carbon dioxide escapes, the amount of carbonic acid will decrease.  

When all is said and done, water and carbon dioxide are two extremely stable molecules.  Since
$$\Delta G = -RT\ln K$$
Their stability is what drives the equilibrium to the left and why carbonic acid appears to be (relatively) so unstable.