A sample of N2O3(g) has a pressure of 0.046 atm . The temperature (in K) is then doubled and the N2O3 undergoes complete decomposition to NO2(g) and NO(g). Find the total pressure of the mixture of gases assuming constant volume and no additional temperature change. Enter your answer numerically, in terms of atm.

So, the sample of N₂O₃ will only have its temperature doubled, with the same volume and the same number of moles. Temperature and pressure are directly related, so if one increases the other also increases, then the pressure must double to 0.092 atm.

The decomposition occurs:

N₂O₃(g) ⇄ NO₂(g) + NO(g)

So, 1 mol of N₂O₃ will produce 2 moles of the products (1 of each), the n will double. The volume and the temperature are now constants, and the pressure is directly proportional to the number of moles, so the pressure will double to 0.184 atm.

niharikadhawanarora niharikadhawanarora · Answer 2 · 2021-10-23T11:51:21+02:00

Formula to be used:

[tex]PV= nRT[/tex] is the ideal gas equation.

Where P denotes pressure, V denotes volume, n is the number of moles, R denotes the gas constant, and T denotes temperature.

As a result, the temperature of the [tex]N_{2} O_{3}[/tex] sample will only be doubled, with the same volume and number of moles. Because temperature and pressure are inextricably linked, if one rises, the other must rise as well, resulting in a pressure of 0.092 atm.
The following are the steps in the breakdown process: Decomposition reaction: [tex]N_{2} O_{3}[/tex] ⇄ [tex]NO_{2}[/tex] + [tex]NO[/tex]
As a result, 1 mole of [tex]N_{2} O_{3}[/tex] produces 2 moles of products (1 of each), and n doubles. The pressure will double to 0.184 atm since the volume and temperature are now constants, and the pressure is directly proportional to the number of moles.

Thus, the total pressure is 0.184 atm.

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