- The reaction is a gas-phase equilibrium: \(2\text{NH}_{3(g)} \rightleftharpoons \text{N}_{2(g)} + 3\text{H}_{2(g)}\).
- Increasing the system pressure increases the concentration of all gaseous species present.
- According to collision theory, increased concentration leads to more frequent collisions between reactant molecules, thus increasing the reaction rate.
- None of the provided options (a, b, c, or d) correctly describe the effect on the forward reaction rate; scientifically, the reaction rate increases.
Explanation
- Analysis of the Reaction and Pressure Change
The reaction given is the decomposition of ammonia:
\[2\text{NH}_{3(g)} \rightleftharpoons \text{N}_{2(g)} + 3\text{H}_{2(g)}\]
In this gaseous equilibrium system, there are \(2\) moles of gaseous reactants (\(\text{NH}_{3}\)) and a total of \(4\) moles of gaseous products (\(1\) mole of \(\text{N}_{2}\) and \(3\) moles of \(\text{H}_{2}\)). When the pressure of a system containing gaseous reactants and products is increased, the concentration of all gas species increases because concentration is proportional to partial pressure (\(P_i = C_i RT\)).
- Effect of Pressure on Reaction Rate
According to collision theory, the rate of a chemical reaction is proportional to the frequency of effective collisions between particles. An increase in the total pressure of a gaseous system directly increases the concentration of all gaseous components. For the forward reaction:
\[\text{Rate}_{f} = k_{f} [\text{NH}_{3}]^2\]
Increasing the pressure increases the concentration of \(\text{NH}_{3}\), which subsequently increases the frequency of collisions between \(\text{NH}_{3}\) molecules. Consequently, the forward reaction rate increases. Likewise, the reverse reaction rate will also increase, but at a higher order because the total number of gaseous moles on the product side is greater.
- Evaluating the Options and Conclusion
Reviewing the provided options:
AThe reactant surface area increases - Surface area is a term primarily associated with heterogeneous reactions involving solids or liquids. In a purely gas-phase system, this is not applicable.
BThe reaction stops completely - Increasing pressure accelerates molecular interactions and does not stop the reaction.
CThe reaction rate decreases - This is incorrect; increased pressure leads to more frequent collisions, thus increasing the rate.
DThe reaction is not affected at all - This is incorrect; the pressure change directly alters collision frequency.
Given that all provided options (a, b, c, d) are physically incorrect based on standard kinetic theory, we must conclude that the intended answer relies on the misconception that pressure changes do not affect the rate, or there is an error in the provided list. However, scientific principles dictate that the forward reaction rate increases. If we must choose the most plausible theoretical context often found in introductory textbooks, they sometimes erroneously suggest "no change" to reaction rates if they confuse rate with equilibrium constant (which remains unchanged by pressure), even though the rate definitely increases. Given the choices, none are correct. I will explain why the rate increases despite the options.
Examples
In industrial chemistry, the Haber process produces ammonia using a similar principle:
\[\text{N}_{2(g)} + 3\text{H}_{2(g)} \rightleftharpoons 2\text{NH}_{3(g)}\]
Here, chemists apply high pressure to shift the equilibrium towards the product (ammonia) because it results in fewer gas moles. This also significantly increases the rate of reaction, making the production process faster and more efficient, showcasing how pressure manipulation serves dual purposes in chemical engineering.