QUESTION SETUP: Assume at 273K and a CO2 pressure of 3.6 atm, the aqueous solubi

Question

QUESTION SETUP: Assume at 273K and a CO2 pressure of 3.6 atm, the aqueous solubility of CO2 is 24.8 ml CO2 per liter. What is the molarity of a saturated water solution when the CO2 is under its normal partial pressure in air of .000395atm.

MY QUESTION: The answer is below, but I'm confused by the first step. Where is the 1/22.4LCO2 derived from? What am I missing? If it is a calculated number, please provide details. Thanks!

Example 5: Henry's Law Application Assume at 273K and a CO2 pressure of 3.6 atm, the aqueous solubility of CO2 is 24.8 ml CO2 per liter. What is the molarity of a saturated water solution when the CO2 is under its normal partial pressure in air of.000395atm. Step 1: We need to calculate Henry's Constant for CO2 when the partial pressure of CO2 is 3.6 atm. To do this, we must first calculate the molarity of CO2 under these conditions: 0.0248Lco, * 22 molCO2 0.0248LCO222.AL0 Molarity = 0.0011|MC02 1Lsolution Step 2: Next we must solve for Henry's Constant under these conditions: MCO2 atm H 0.000308 Step 3: Solve for the concentration (molarity) using Henry's Law

Explanation / Answer

1 mole of any gas at STP( standard temperature & pressure) occupies 22.4 L

That's why we write (1 mol CO2 )/ (22.4 L CO2)

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