Question

Carbonated Beverages Use Le Châtelier's principle to explain how a shift in the equilibrium \(\mathrm{H}_{2} \mathrm{CO}_{3}(\mathrm{aq}) \rightleftharpoons\) \(\mathrm{H}_{2} \mathrm{O}(1)+\mathrm{CO}_{2}(\mathrm{g})\) causes a soft drink to go flat when its container is left open.

Short answer

When a container of a carbonated beverage is opened, the pressure inside decreases, causing the equilibrium of the reaction \(\mathrm{H}_{2} \mathrm{CO}_{3}(\mathrm{aq}) \rightleftharpoons \mathrm{H}_{2} \mathrm{O}(1)+\mathrm{CO}_{2}(\mathrm{g})\) to shift. According to Le Châtelier's Principle, the equilibrium shifts to the right, producing more \(\mathrm{CO}_{2}(\mathrm{g})\). This gas is released into the atmosphere, decreasing the concentration of dissolved \(\mathrm{CO}_{2}\) in the soft drink, leading to reduced carbonation and causing the drink to go flat.

Step-by-step solution

Understand Le Châtelier's Principle

Le Châtelier's principle states that when a system at equilibrium is subjected to a change in concentration, pressure, or temperature, the equilibrium will shift to counteract the change and partially re-establish equilibrium. In this case, we will consider how the change in pressure (due to the container being opened) affects the equilibrium between \(\mathrm{H}_{2} \mathrm{CO}_{3}\), \(\mathrm{H}_{2} \mathrm{O}\), and \(\mathrm{CO}_{2}\).

Identify the Change when the Container is Opened

When a container of a carbonated beverage is opened, the pressure inside the container decreases as the container is exposed to atmospheric pressure. This change in pressure affects the equilibrium between \(\mathrm{H}_{2} \mathrm{CO}_{3}(\mathrm{aq})\), \(\mathrm{H}_{2} \mathrm{O}(1)\), and \(\mathrm{CO}_{2}(\mathrm{g})\).

Applying Le Châtelier's Principle

Since there is a decrease in pressure, according to Le Châtelier's Principle, the equilibrium will shift toward the side of the reaction that has more moles of gas in order to counteract the decrease in pressure. In this case, the equilibrium will shift to the right, towards the side with \(\mathrm{CO}_{2}(\mathrm{g})\). The reaction will now proceed as follows: \(\mathrm{H}_{2} \mathrm{CO}_{3}(\mathrm{aq}) \rightarrow \mathrm{H}_{2} \mathrm{O}(1)+\mathrm{CO}_{2}(\mathrm{g})\)

Understanding the Effect on Carbonation

As the equilibrium shifts to the right, more \(\mathrm{CO}_{2}\) gas is produced. Since the container is open, the \(\mathrm{CO}_{2}\) gas is released into the atmosphere, resulting in a decrease in the concentration of dissolved \(\mathrm{CO}_{2}\) in the soft drink. This leads to a decrease in carbonation (fizz) and eventually causes the soft drink to go flat.

Key concepts

Chemical Equilibrium

Chemical equilibrium is a state in chemical reactions where the rate of the forward reaction equals the rate of the reverse reaction, leading to a constant concentration of the reactants and products. It is central to understanding numerous processes, including the carbonation of beverages. In equilibrium, any change can push the system to adjust in order to maintain this balance, a concept described by Le Châtelier's principle.

For example, in the equilibrium between carbonic acid, water, and carbon dioxide (\(\mathrm{H}_{2}\mathrm{CO}_{3}(\mathrm{aq}) \rightleftharpoons \mathrm{H}_{2}\mathrm{O}(l) + \mathrm{CO}_{2}(\mathrm{g})\)), alterations in conditions such as pressure trigger a response in the system to restore the equilibrium. When these changes occur, we can predict the direction of the shift—whether it's toward more reactants or products—using Le Châtelier's principle.

Carbonation of Beverages

The fizzy sensation in carbonated beverages like sodas comes from the dissolution of carbon dioxide gas under high pressure. When CO₂ dissolves in water, it forms carbonic acid (\(\mathrm{H}_{2}\mathrm{CO}_{3}\)), contributing to the tangy flavor and carbonation.

Chemical Dynamics in a Closed Container

In a sealed container, a kind of chemical equilibrium is established between dissolved carbon dioxide and carbonic acid. The pressure inside is greater than the atmospheric pressure, maintaining high levels of dissolved CO₂. If this equilibrium gets disturbed, for instance, when a bottle is shaken, the balance is promptly restored, keeping the drink carbonated as long as the container remains sealed.

Effects of Pressure on Equilibrium

The impact of changing pressure on chemical equilibrium is articulated by Le Châtelier's principle. Increases in pressure will favor the side of an equilibrium with fewer moles of gas, while decreases will favor the side with more moles.

Opening the Beverage

When you open a carbonated beverage, the pressure inside the container drops substantially, interrupting the established equilibrium. Following Le Châtelier's principle, the shift in pressure causes the equilibrium to adjust by favoring the production of more gas—forming more carbon dioxide from carbonic acid and water.

As the system seeks a new equilibrium at the lower pressure, it accelerates the release of CO₂ gas until the pressures inside and outside the container become equal. This increased production of gaseous CO₂ is what leaves the beverage flat, as the pressure decrease drives the reaction to produce more gas which then escapes into the atmosphere.