Gas Laws

At STP, 1 mole of any ideal gas occupies 22.4 L. Moles = 11.2 / 22.4 = 0.5 mol. Molar mass = 13 / 0.5 = 26 g/mol. C H has a molar mass of 26 g/mol.

According to Graham’s Law, lighter gases (H ) effuse faster than heavier ones (O ). More hydrogen escapes, reducing its partial pressure more than oxygen. Hence, oxygen remains in higher proportion in A.

At constant temperature, the pressure of a gas decreases as the volume increases, as described by Boyle's Law (P 1/V).
The other statements are correct: increasing temperature increases the pressure at constant volume, increasing volume increases the volume at constant pressure, and the Kelvin scale is preferred in gas law calculations as it avoids negative temperatures.

We are given that 100 g of both O and He are in separate containers of equal volume at the same temperature.
According to the ideal gas law , the pressure depends on the number of moles of the gas. The number of moles of O and He can be calculated as: Thus, the moles of He are significantly greater than the moles of O .
Since the gases are in the same volume and at the same temperature, the pressure is directly proportional to the number of moles.
Therefore, the pressure of He will be greater than that of O .
Thus, the correct answer is (d).

The Haber process is an equilibrium reaction, where , indicating that the reaction is exothermic. According to Le Chatelier's Principle, increasing the pressure favours the side of the reaction with fewer moles of gas, which in this case is the production of ammonia. The balanced reaction is: Here, 4 moles of reactants (N2 and H2) are converted into 2 moles of products (NH3), so increasing the pressure shifts the equilibrium towards the production of ammonia, thereby increasing the yield. Additionally, a higher temperature increases the rate of the reaction but decreases the yield, as the forward reaction is exothermic. If the temperature is too high, the reverse reaction becomes more favored. Therefore, a compromise temperature is used in practice, which provides a balance between a reasonable rate of reaction and an acceptable yield. Thus, the correct statement is (d) A higher pressure would ensure a higher yield at a higher cost, as higher pressure favours ammonia production but requires more energy and equipment to sustain.

We use the Ideal Gas Law to calculate the molecular mass: Where: Using given values for P, V, and T, and applying the equation: Substitute the known values:

First, we use the balanced chemical equation: This shows that for every 4 moles of NH3, 4 moles of NO2 are produced. Since we start with 8 moles of NH3, we will produce 8 moles of NO2. The total initial moles of gas in the container are . The initial pressure is 11 atm, so the total pressure is proportional to the total moles of gas. The final pressure is proportional to the moles of NO2 formed, which is 8 moles. Thus, the partial pressure of NO2 is: Therefore, the correct answer is (a) 4 atm.