In water saturated air, the mole fraction of water vapour is $0.02$. If the total pressure of the saturated air is $1.2\, atm$, the partial pressure of dry air is :

Question

Which one of the following is not correct for an ideal solution?

Answer 1

To find the partial pressure of dry air in a mixture where the total pressure is given and the mole fraction of water vapor is known, we can use Dalton's Law of Partial Pressures.

According to Dalton's Law, the total pressure of a mixture of gases is equal to the sum of the partial pressures of the individual gases:

P_{\text{total}} = P_{\text{dry air}} + P_{\text{water vapor}}

We are given:

Total pressure, P_{\text{total}} = 1.2 \, \text{atm}
Mole fraction of water vapor, x_{\text{water vapor}} = 0.02

The partial pressure of a component in a gas mixture can be calculated using its mole fraction and the total pressure:

P_{\text{water vapor}} = x_{\text{water vapor}} \times P_{\text{total}}

Substituting the known values:

P_{\text{water vapor}} = 0.02 \times 1.2 \, \text{atm} = 0.024 \, \text{atm}

Now, using Dalton's Law, we can find the partial pressure of dry air:

P_{\text{dry air}} = P_{\text{total}} - P_{\text{water vapor}}

Substitute the values:

P_{\text{dry air}} = 1.2 \, \text{atm} - 0.024 \, \text{atm} = 1.176 \, \text{atm}

Therefore, the partial pressure of dry air is 1.176 \, \text{atm}.

Hence, the correct answer is 1.176 atm.

Answer 2