N2H4 is decomposed in a closed container according to this reaction: N2H4 → N2 + 2H2 . If the temperature is constant and the initial pressure of N2H4 is 800mm, what will the pressure in millimeters be after the decomposition to N2 + 2H2? Im having trouble with this problem because I don't see how can I find the temperature , wich is constant.
You don't need the temperature but you do need Keq unless, as I suspect, the author of the problem wants us to assume that the decomposition is COMPLETE. If that's the case, it is done this way.
.......N2H4 ==> N2 + 2H2
I......800......0.....0
C.....-800......800...1600
E.......0.......800...1600
So total pressure of the system will be 800 + 1600 = ? mm Hg.
To find the pressure of the products (N2 and H2) after the decomposition of N2H4, you don't need to know the actual temperature; you only need to know that the temperature is constant, which means it remains the same throughout the reaction.
The given reaction, N2H4 → N2 + 2H2, shows that one mole of N2H4 decomposes to produce one mole of N2 and two moles of H2. Since the mole ratio of N2 to N2H4 is 1:1 and the mole ratio of H2 to N2H4 is 2:1, the volume ratio will also be 1:1 for N2 and 2:1 for H2.
Using the ideal gas law (PV = nRT), where P is the pressure, V is the volume, n is the number of moles, R is the ideal gas constant, and T is the temperature, we can relate the initial pressure, volume, and the number of moles of N2H4 to the final pressure of the products.
Since the volume is constant and the temperature is constant, we can simplify the equation to PV = constant.
Initially, the moles of N2H4 are given by the ideal gas law equation:
n1 = (P1V1) / RT
After decomposition, one mole of N2H4 produces one mole of N2, so the number of moles of N2 is also equal to n1.
Therefore, the final pressure, P2, can be calculated using the number of moles of N2 as follows:
P2 = (n2RT) / V
Since we have the initial pressure (P1), the initial volume (V1), and the number of moles of N2H4 (n1), we can calculate the final pressure (P2) by substituting these values into the equation.
Here are the steps to solve for the final pressure:
Step 1: Convert the initial pressure from millimeters (mm) to standard atmospheric pressure units (atm). 1 atm = 760 mm.
P1 = 800 mm / 760 mm/atm = 1.05 atm
Step 2: Substitute the known values into the equation for n1.
n1 = (P1V1) / RT
Step 3: Use the stoichiometry of the reaction to find the number of moles of N2.
n2 = n1
Step 4: Substitute the known values into the equation for P2 to find the final pressure.
P2 = (n2RT) / V
Note: Make sure to use consistent units in the calculations (e.g., convert volume to liters, temperature to Kelvin, and pressure to atmospheres).
By following these steps, you can calculate the final pressure (P2) in millimeters.