At 500^∘ C, K for the formation of ammonia from nitrogen and...

At $500^{\circ} \mathrm{C}, K$ for the formation of ammonia from nitrogen and hydrogen gases is $1.5 \times 10^{-5}$ $$ \mathrm{N}_{2}(g)+3 \mathrm{H}_{2}(g) \rightleftharpoons 2 \mathrm{NH}_{3}(g) $$ Calculate the equilibrium partial pressure of hydrogen if the equilibrium partial pressures of ammonia and nitrogen are 0.015 atm and 1.2 atm, respectively.

At $500^{\circ} \mathrm{C}, K$ for the formation of ammonia from nitrogen and hydrogen gases is $1.5 \times 10^{-5}$
$$
\mathrm{N}_{2}(g)+3 \mathrm{H}_{2}(g) \rightleftharpoons 2 \mathrm{NH}_{3}(g)
$$
Calculate the equilibrium partial pressure of hydrogen if the equilibrium partial pressures of ammonia and nitrogen are 0.015 atm and 1.2 atm, respectively.

Key Concepts

Reaction Stoichiometry

Reaction stoichiometry involves the quantitative relationships between reactants and products in a chemical reaction. It is used to determine the exponent applied to each partial pressure in the equilibrium constant expression based on the balanced chemical equation, thereby connecting the physical quantities in the reaction to the mathematical formulation of equilibrium.

Partial Pressure

Partial pressure is the pressure exerted by an individual gas in a mixture of gases. In the context of equilibrium calculations, knowing the partial pressures allows us to determine the reaction’s status and to calculate unknown pressures using the equilibrium constant expression. This concept is critical in the study of gas-phase reactions according to Dalton’s Law.

Equilibrium Constant (Kp)

The equilibrium constant in terms of partial pressures (Kp) is a dimensionless number that relates the products and reactants’ pressures at equilibrium, based on their stoichiometric coefficients. It is defined by the ratio of the partial pressures of products to those of reactants, each raised to the power of their respective coefficients, and it quantifies the position of equilibrium for a gaseous reaction.

Chemical Equilibrium

Chemical equilibrium occurs when the forward and reverse reactions proceed at the same rate so that the concentrations or partial pressures of reactants and products remain constant over time. This concept is fundamental in understanding how reactions reach a balanced state without any net change, even though the reactions continue to occur at the molecular level.

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