Question

In a spontaneous acid-base reaction in the gas phase:

NH₃​(g)+HCl(g)→NH₄​Cl(s)

During an experiment, 11.00 g of solid NH₄​Cl was added to a 500.0 mL flask at 25.0°C. Given that the equilibrium constant (Kc​) for this reaction at this temperature is 8.9×10⁻¹, determine the equilibrium partial pressures (in atm) of NH₃​ and HCl.

Answer 1

The calculation of equilibrium partial pressures of NH₃ and HCl from solid NH₄Cl cannot be completed with the given data, as the degree of dissociation or initial pressures of the gases is not provided.

Step-by-step explanation:

The student's question pertains to determining the equilibrium partial pressures of NH3 and HCl gases in a reaction involving ammonium chloride in the gas phase. To find these pressures, we must consider the stoichiometry of the balanced chemical reaction NH3(g) + HCl(g) → NH4Cl(s) and the provided equilibrium constant (Kc) at a specified temperature.

When 11.00 g of solid NH4Cl is added to a 500.0 mL flask, the substance can dissociate into NH3 and HCl gases until the equilibrium is reached. The equilibrium constant given does not directly apply to this situation as it pertains to different conditions. Calculating the partial pressures would require knowledge of the degree of dissociation, which can't be determined with the given data.

In the thermal dissociation of NH4Cl, as mentioned in provided references, the equilibrium constant Kp or Kc can be used if the partial pressures or concentrations of the resultant gases are known at equilibrium. However, without knowing the degree of dissociation or initial pressures, we cannot proceed with the calculation. In instances where pressures or concentrations at equilibrium are given, one could apply the expression for Kp or Kc to solve for the unknown.