Aqueous solubility of CO2 is 24.8 ml of CO2 per liter at 273K and a CO2 partial pressure of 3.6 atm... find molarity

Question

Aqueous solubility of CO2 is 24.8 ml of CO2 per liter at 273K and a CO2 partial pressure of 3.6 atm. What is the molarity of a saturated water solution when the CO2 is under its normal partial pressure in air of 3.95x10-4 atm?

Answer 1

You first need to determine Henry's constant (Kh) at this temperature:

 

Kh = P(CO2)/[CO2(aq)]

 

where P(CO2) is the partial pressure of CO2 and [CO2(aq)] is the concentration of CO2 in solution. Thus,

 

Kh = (3.6 atm)/[(24.8 mL CO2)/(1000 mL H2O)] = 145.2 atm*mL H2O/mL CO2

 

Now use the new pressure and the calculated Henry's constant at 273K to solve for the concentration of CO2:

 

[CO2(aq)] = P(CO2)/Kh = (3.95x10-4 atm)/(145.2 atm*mL H2O/mL CO2) = 2.72x10-6 mL CO2/mL H2O

 

At this point you can probably assume that the volume of H2O won't change much after gas dissolution and looking up the density of CO2 at 273K (1.964 g/L) we can determine the molarity of CO2 in solution:

 

M = (2.72x10-6 mL CO2/mL H2O) x [(1.964 g CO2)/(1000 mL CO2)] x [(1 mole CO2)/(44.01 g CO2)] x [(1000 mL H2O)/(1 L H2O)] = 1.21x10^-7 M or 0.121 uM.