🤓 Based on our data, we think this question is relevant for Professor Co's class at OSU.

The partial pressure of O_{2} in air at sea level is 0.21 atm . The solubility of O_{2} in water at 20 ^{o}C , at 1 atm gas pressure, is 1.38 10^{-3} M (from Table 13.1 in your textbook).

Using Henrys law and the data in the introduction, calculate the molar concentration of O_{2} in the surface water of a mountain lake saturated with air at 20 ^{o}C and an atmospheric pressure of 670 torr .

We’re being asked to** calculate the molar concentration of O**_{2} in the surface water of a mountain lake **saturated with air at 20 **^{o}**C** and an** ****atmospheric pressure of 670 torr**.

Recall that the solubility of a gas is given by ** Henry’s law**:

$\overline{){{\mathbf{S}}}_{{\mathbf{gas}}}{\mathbf{=}}{{\mathbf{k}}}_{{\mathbf{H}}}{\mathbf{\xb7}}{{\mathbf{P}}}_{{\mathbf{gas}}}}$

where:

**S _{gas}** = solubility of the gas (in mol/L or M)

**k _{H}** = Henry’s law constant for the gas

**P _{gas}** = partial pressure of the gas

**For this problem, we need to do the following steps:**

**Step 1***:** Calculate the Henry’s Law constant, k_{H }of O_{2} based on the given solubility and pressure at 20 ^{o}C using Henry's Law*

**Step 2***: **Calculate the mole fraction of O _{2} in the air at the mountain lake *

**Step 3****: ***Determine the partial pressure of O _{2} at the mountain lake from the given atmospheric pressure *

*Ste**p 4**:*Calculate the molar concentration or solubility* of O _{2} in the surface water *

Henry's Law