A gaseous mixture of 80.0g of oxygen and 56.0 of nitrogen has a total pressure

Explanation

To answer this question, we need to apply Dalton's Law of Partial Pressures, which states that the total pressure of a gas mixture is the sum of the partial pressures of the individual gases.

First, let's find the mole fraction of oxygen and nitrogen in the mixture. We can do this by dividing the given mass of each gas by its molar mass.

Oxygen: \( \frac{80.0 \, g}{32.0 \, g/mol} = 2.5 \, mol \)

Nitrogen: \( \frac{56.0 \, g}{28.0 \, g/mol} = 2.0 \, mol \)

Now, we need to calculate the mole fraction of each gas. The mole fraction is the ratio of the moles of a particular gas to the total moles of the gas mixture.

Mole fraction of oxygen: \( \frac{2.5 \, mol}{2.5 \, mol + 2.0 \, mol} = \frac{2.5}{4.5} \)

Mole fraction of nitrogen: \( \frac{2.0 \, mol}{2.5 \, mol + 2.0 \, mol} = \frac{2.0}{4.5} \)

Now that we have the mole fractions, we can use Dalton's Law of Partial Pressures to find the partial pressure of oxygen.

\( P_{O_2} = X_{O_2} \times P_{total} \)

\( P_{O_2} = \frac{2.5}{4.5} \times 1.8 \, atm \)

\( P_{O_2} = 1.0 \, atm \)

So, the partial pressure of oxygen in the mixture is1.0 atm, which corresponds toOption B.