For this problem, we need to determine the **partial pressure of Oxygen**, in a mixture of **2.0 g O _{2} **for every

Recall that based on Dalton’s Law, the **partial pressure** of each gas is the **total pressure**** ****multiplied by the mole fraction of that gas.**

$\overline{){\mathbf{}}{{\mathbf{P}}}_{{\mathbf{gas}}}{\mathbf{=}}{{\mathbf{X}}}_{{\mathbf{gas}}}{\mathbf{\xb7}}{{\mathbf{P}}}_{{\mathbf{Tot}}}\mathbf{}}$

Where:

*P _{Tot} = total Pressure in atm*

*P° = partial pressure of a gas in atm *

*X = mole fraction of gas*

We will need to do these steps to **solve the partial pressure of O _{2}:**

1. * Solve for # of moles of ***O _{2 }**and

Do a **mass to mole calculation** using the molar mass

2. Find the **mole fraction of O _{2 }**using the equation:

$\overline{){{\mathbf{X}}}_{{\mathbf{O}}_{\mathbf{2}}}{\mathbf{}}{\mathbf{=}}\frac{\mathbf{moles}\mathbf{}{\mathbf{O}}_{\mathbf{2}}}{\mathbf{moles}\mathbf{}{\mathbf{O}}_{\mathbf{2}}\mathbf{}\mathbf{+}\mathbf{}\mathbf{moles}\mathbf{}\mathbf{He}}}$

3. Solve for partial pressure of O_{2} given a **total pressure of 8.2 atm. **

A heliox deep-sea diving mixture contains 2.0 g of oxygen to every 98.0 g of helium. What is the partial pressure of oxygen when this mixture is delivered at a total pressure of 8.2 atm?

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