Question

Baking soda (sodium bicarbonate) undergoes thermal decomposition as follows: $$ 2 \mathrm{NaHCO}_{3}(s) \rightleftarrows \mathrm{Na}_{2} \mathrm{CO}_{3}(s)+\mathrm{CO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(g) $$ Would we obtain more \(\mathrm{CO}_{2}\) and \(\mathrm{H}_{2} \mathrm{O}\) by adding extra baking soda to the reaction mixture in (a) a closed vessel or (b) an open vessel?

Short answer

More \(\mathrm{CO}_{2}\) and \(\mathrm{H}_{2} \mathrm{O}\) are obtained in an open vessel.

Step-by-step solution

Understanding the Reaction

The reaction in question is the thermal decomposition of baking soda, which is: \[ 2 \mathrm{NaHCO}_{3}(s) \rightleftarrows \mathrm{Na}_{2}\mathrm{CO}_{3}(s)+\mathrm{CO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(g) \] This reaction is reversible, and it involves solid sodium bicarbonate (\mathrm{NaHCO}_{3}) decomposing into sodium carbonate (\mathrm{Na}_{2} \mathrm{CO}_{3}), carbon dioxide (\mathrm{CO}_{2}), and water vapor (\mathrm{H}_{2} \mathrm{O}).

Applying Le Chatelier's Principle

To determine whether the addition of more baking soda will yield more \(\mathrm{CO}_{2}\) and \(\mathrm{H}_{2} \mathrm{O}\), consider Le Chatelier's Principle. Increasing the concentration of a reactant typically shifts the equilibrium towards the products in order to minimize the change, which would increase the amount of \(\mathrm{CO}_{2}\) and \(\mathrm{H}_{2} \mathrm{O}\) produced.

Case Analysis - Closed Vessel

In a closed vessel, the system is sealed, meaning that the gaseous products (\(\mathrm{CO}_{2}\) and \(\mathrm{H}_{2} \mathrm{O}\)) cannot escape. As \(\mathrm{NaHCO}_{3}\) is added, the reaction will shift to the right to produce more \(\mathrm{CO}_{2}\) and \(\mathrm{H}_{2} \mathrm{O}\). However, as these gases build up, equilibrium is reached again without releasing excess gas, which can limit their maximum concentration.

Case Analysis - Open Vessel

In an open vessel, gases like \(\mathrm{CO}_{2}\) and \(\mathrm{H}_{2} \mathrm{O}\) can escape the system as they are formed. Adding more \(\mathrm{NaHCO}_{3}\) will continuously shift the reaction towards the products, increasing \(\mathrm{CO}_{2}\) and \(\mathrm{H}_{2} \mathrm{O}\) production as these gases escape, preventing the reaction from reaching equilibrium and allowing more gas to be continually produced.

Conclusion

Comparing the two scenarios, more \(\mathrm{CO}_{2}\) and \(\mathrm{H}_{2} \mathrm{O}\) are produced in an open vessel because the gases can escape, allowing the reaction to keep shifting right without limitation.

Key concepts

Sodium Bicarbonate Decomposition

In chemistry, sodium bicarbonate, or baking soda, is a common compound with various uses, including in cooking and as a cleaning agent. A fascinating phenomenon occurs when sodium bicarbonate is subjected to heat; it undergoes thermal decomposition. This process involves the breakdown of sodium bicarbonate (\(\mathrm{NaHCO}_{3}\)) into three simpler products: sodium carbonate (\(\mathrm{Na}_{2}\mathrm{CO}_{3}\)), carbon dioxide (\(\mathrm{CO}_{2}\)), and water vapor (\(\mathrm{H}_{2}\mathrm{O}\)).This chemical reaction is important because it is a reversible reaction, meaning it can proceed in both the forward and backward directions. When heated, the forward reaction dominates, producing the solid and gaseous products. The reaction can be represented by the equation:\[2 \mathrm{NaHCO}_{3}(s) \rightleftarrows \mathrm{Na}_{2} \mathrm{CO}_{3}(s)+\mathrm{CO}_{2}(g)+\mathrm{H}_{2} \mathrm{O}(g)\]Understanding this decomposition is crucial for applications like baking, where the release of carbon dioxide helps dough rise. Thermal decomposition is not only useful in baking but also provides insights into chemical behavior and reaction dynamics.

Le Chatelier's Principle

Le Chatelier's Principle is a fundamental concept in chemistry that helps us predict the effect of changes on a chemical system at equilibrium. It states that if a dynamic equilibrium is disturbed by changing the conditions, the position of equilibrium shifts to counteract the change and re-establish equilibrium.In the context of sodium bicarbonate decomposition, if extra baking soda is added to the mixture, Le Chatelier's Principle can be applied. Adding more sodium bicarbonate increases the concentration of reactants and disturbs the equilibrium. According to the principle, the equilibrium will shift towards the production of more products (\(\mathrm{CO}_{2}\) and \(\mathrm{H}_{2}\mathrm{O}\)) to balance the change.This shift in equilibrium is particularly evident in closed and open systems:

• In a closed system, gases cannot escape, so the shift eventually reaches a new equilibrium.
• In an open system, gases like \(\mathrm{CO}_{2}\) and \(\mathrm{H}_{2}\mathrm{O}\) can escape, continually driving the reaction forward.

Understanding Le Chatelier's Principle allows chemists and students to predict and control the outcomes of reactions, making it a powerful tool in both academic and practical settings.

Equilibrium Reactions

Equilibrium reactions are a fundamental aspect of chemistry, where the rate of the forward reaction equals the rate of the reverse reaction, resulting in a stable balance between products and reactants. In the decomposition of sodium bicarbonate, equilibrium is achieved when the decomposition rate into sodium carbonate, carbon dioxide, and water vapor matches their recombination.The concept of equilibrium is dynamic rather than static, meaning that reactions continue to occur, but their effects cancel out, maintaining a stable concentration of substances. Several factors can affect equilibrium, including pressure, temperature, and concentration.In the given reaction, the presence of gaseous products (\(\mathrm{CO}_{2}\) and \(\mathrm{H}_{2}\mathrm{O}\)) allows us to explore how changes affect equilibrium:

• Increasing pressure or temperature can shift equilibrium.
• Removing products (especially gases) in an open system shifts equilibrium to favor product formation continuously.

Understanding the nature of equilibrium reactions helps in manipulating chemical processes, making it essential knowledge in chemistry. Mastery of these concepts enables the prediction and optimization of chemical yields in both laboratory and industrial scenarios.