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The decomposition of nitrogen dioxide in the gas phase Occurs according to the following equation at 50 °C: \mathrm{NO}_{2}(\mathrm{~g}) \rightarrow 1 / 2 \mathrm{O}_{2}(\mathrm{~g})+\mathrm{NO}(\mathrm{g}) \text { The rate law for

the reaction has been found to be } R=0.543 \mathrm{M}^{-1} \mathrm{~s}^{-1}\left[\mathrm{NO}_{2}\right]^{2} \text { at } 50^{\circ} \mathrm{C} \text { . } a. Given an initial concentration of 0.0240 M of NO2, what would the half-life for the reaction be? b. How much NO2 would be left after 50.0 seconds if the starting concentration of NO2 was 0.0240 M? c. Find the relationship between the total pressure, the fraction of molecules that have reacted and the initial pressure. d. If the initial concentration of NO2 is 0.0240 M, what would be the total pressure in units of atm at the start of the reaction, after 50.0 seconds and at the end of the reaction? Assume all species behave as ideal gasses.

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