Clarke Mckenzie
06/06/2024 · High School
Consider the following equilibrium system: \[ \mathrm{PCl}_{5}(\mathrm{~g}) \rightleftharpoons \mathrm{PCl}_{3}(\mathrm{~g})+\mathrm{Cl}_{2}(\mathrm{~g}) \] A 10.00 L evacuated flask is filled with 0.4414 mol \( \mathrm{PCl}_{5}(\mathrm{~g}) \) at \( \mathbf{2 9 8 . 3} \mathrm{K} \). The temperature is then raised to 503.3 K , where the decomposition of \( \mathrm{PCl}_{5} \) gas takes place to an appreciable extent. When equilibrium is established, the total pressure in the flask is \( \mathbf{2 . 5 9 7} \) atm.
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The equilibrium concentrations of \( \mathrm{PCl}_{5} \), \( \mathrm{PCl}_{3} \), and \( \mathrm{Cl}_{2} \) are determined by solving for \( x \), the amount of \( \mathrm{PCl}_{5} \) that decomposes. The equilibrium constant \( K_p \) is calculated using the partial pressures of the gases.
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