Final answer:
The question pertains to determining the equilibrium concentrations in the reversible reaction H2(g) + I2(g) ⇌ 2HI(g), using the equilibrium constant K and initial concentrations to solve for the unknown values.
Step-by-step explanation:
The question concerns the chemical equilibrium of the reaction involving hydrogen gas (H₂) and iodine gas (I₂) to form hydrogen iodide (HI). In chemical equilibrium, the rates of the forward and reverse reactions are equal, meaning that the concentrations of reactants and products remain constant over time. The reaction can be represented as H₂(g) + I₂(g) ⇌ 2HI(g), where the double arrows indicate that the reaction is reversible.
To solve for equilibrium concentrations, the equilibrium constant, K, is used. K is expressed as a ratio that involves the concentrations of the products raised to the power of their coefficients in the balanced equation, divided by the concentrations of the reactants raised to the power of their coefficients. Given an equilibrium constant (K) and initial concentrations of reactants, the equilibrium concentrations of all species participating in the reaction can be calculated. This involves setting up an ICE table (Initial, Change, Equilibrium), applying the equilibrium expression, and solving for the unknown concentrations.
For example, if the initial concentrations and the equilibrium constant are known, as in the provided scenario where K is 50.5, one can use the equilibrium expression in the form of K = [HI]^2 / ([H₂][I₂]) to find the missing equilibrium concentrations.