PRESSURE:
CO2 (g) <----> CO2 (aq) (Δ < 0) - (1)
CO2 (aq) + H20 (l) <----> H2CO3 (aq) - (2)
If pressure is increased, the equilibrium will shift to favour the forward reaction which reduces the pressure, causing the formation of more aqeous CO2 (aq), which means that CO2 (aq) solubility is increased within the soft drink bottle.
Increased CO2 (aq) solubility promotes another forward reaction increasing H2CO3 and thus the acidity of the solution increases.
When the the lid of the bottle is opened, the volume available for the particles within the bottle increases, decreasing the overall pressure. In accordance with LCP, this decrease in pressure prompts the reverse reaction in reaction (1), leading to the formation of more CO2 (g), which is seen through bubbles. This continuous evolution of the CO2 (g) causes the soft drink to eventually become flat/have no CO2 (aq) in it.
TEMPERATURE:
CO2 (g) <----> CO2 (aq) (Δ < 0) - (1)
As shown in equation 1, since Δ < 0, the forward reaction will be exothermic and the reverse reaction will be endothermic.
So if the solution is heated (temperature increased), then the equilibrium would favour the backward reaction which is endothermic, to cool the system by partially counteracting the imposed temperature increase and to thus re-establish equilibrium according to LCP.
Thus, if the temperature is decreased, the system will favour the forward exothermic reaction to heat the system and counteract any cooling that has occurred through the reverse endothermic reaction, re-establishing equilibrium in accordance with LCP.
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So basically, think of LCP as a chain reaction --> once one thing moves forward/to the right, it will keep shifting in the same direction unless another change is imposed upon the closed system, and then the system will re-establish equilibrium by moving in the opposite direction.
