# A mixture of 1.24 g H2 and 1.51 g He is placed in a 1.00 L container at 28°C. How do you calculate the partial pressure of each gas and the total pressure?

Aug 15, 2016

$\text{We follow Dalton's Law of Partial Pressures........}$ to give ${P}_{\text{Total"=P_"dihydrogen"+P_"helium}} = 24.5 \cdot a t m .$

#### Explanation:

$\text{Dalton's Law of Partial Pressures states:..........}$

In a gaseous mixture the pressure exerted by a component gas is the same as the pressure it would exert if it alone occupied the container.

The total pressure is the sum of the partial pressures.

Thus ${P}_{\text{Total}}$ $=$ ${P}_{\text{dihydrogen"+P_"helium}}$.

Assuming ideality, ${P}_{\text{dihydrogen}}$ $=$ $\frac{n R T}{V}$ $=$ $\frac{1.24 \cdot g}{2.016 \cdot g \cdot m o {l}^{-} 1} \times 0.0821 \cdot L \cdot a t m \cdot {K}^{-} 1 \cdot m o {l}^{-} 1 \times 301 \cdot K \times \frac{1}{1.00 \cdot L} = 15.2 \cdot a t m$

AND

${P}_{\text{helium}}$ $=$ $\frac{n R T}{V}$ $=$ $\frac{1.51 \cdot g}{4.003 \cdot g \cdot m o {l}^{-} 1} \times 0.0821 \cdot L \cdot a t m \cdot {K}^{-} 1 \cdot m o {l}^{-} 1 \times 301 \cdot K \times \frac{1}{1.00 \cdot L} = 9.3 \cdot a t m$

Thus ${P}_{\text{Total}}$ $=$ $\left(15.2 + 9.3\right) \cdot a t m$ $=$ $24.5 \cdot a t m$