Open Question
Use Henry’s Law to determine the molar solubility of helium at a pressure of 1.0 atm and 25 °C.
14. Solutions
Henry's Law Calculations
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The atmospheric pressure in a lab is calculated as 1.3 atm. If oxygen gas contributes 62% of this atmospheric pressure, determine its mass (in g) dissolved at room temperature in 25 L of water. The Henry's Law Constant for oxygen in water at this temperature is 5.3 × 10–5 M/atm.
A
1.4 × 10–3 g
B
6.9 × 10–5 g
C
0.055 g
D
0.034 g
Verified step by step guidance1
Calculate the partial pressure of oxygen gas using the given percentage: \( P_{O_2} = 0.62 \times 1.3 \text{ atm} \).
Apply Henry's Law to find the concentration of dissolved oxygen: \( C = k_H \times P_{O_2} \), where \( k_H = 5.3 \times 10^{-5} \text{ M/atm} \).
Substitute the values into Henry's Law equation to find the molarity of dissolved oxygen: \( C = 5.3 \times 10^{-5} \text{ M/atm} \times P_{O_2} \).
Calculate the number of moles of oxygen in 25 L of water using the molarity: \( ext{moles of } O_2 = C \times 25 \text{ L} \).
Convert the moles of oxygen to grams using the molar mass of oxygen (32.00 g/mol): \( ext{mass of } O_2 = ext{moles of } O_2 \times 32.00 \text{ g/mol} \).
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