Finding the Concentration of H3CO2 in a CO2+H2O reaction.

[miniradman]

1. Find the equlibrium constant if there is 0.030moles/L of Carbon Dioxide going in the forward direction and water is in excess to form Carbonic Acid (H2CO3)

Homework Equations

[products]/[reactants] = K

The Attempt at a Solution

I'm trying to figure out how Carbon Dioxide in the atmosphere will affect the chemistry of the oceans. From my research I've found that CO2_(g) = CO2_(aq) in water is 1.45 g/L where I deduced a equilbrium of 0.030 moles/L. My next step is where I have to figure out how much Carbonic Acid is made (CO2 + H₂0 = H₂CO₃) when I know that the equlibrium for the forward direction is 0.039 L/mol. So:

[H₂CO₃]/[CO₂][H₂O]

which based on what I know translates to
[H₂CO₃]/(0.03)[[H₂O] = 0.039L/mol

However, L/mol is not the same as mol/L and I need to find the concentration of [H₂CO₃] because after that I need to figure out the K values for H₂CO₃ = H⁺+HCO₃
and
HCO₃ = H⁺+CO₃
equilibriums

*sorry, I do apologize but I'm not sure how to get the "reversible reaction" symbol on the physics forums, so I just used an "=" sign *

ps. I was read on wikipedia that the solubility of Carbon Dioxide is 1.45g/L at 25 degrees C at 100kpa. The pressure, is that including the rest of the atmosphere? or just the carbon dioxide alone? because I've calculate the partial pressure of the CO₂to be 11.6139 pa at sea level.

 

[miniradman]

Should I give anymore detail?
More explanation?
More of my working Out?
was the layout incorrect?
... or do people just don't like me...

 

[Borek]

There are thick books written on the subject. Just carbonic acid is not enough to analyze anything, although you can be sure increasing partial pressure of CO₂ lowers water pH. Google for seawater chemistry.

 

[chemisttree]

Yes, it's a http://aslo.org/lo/toc/vol_18/issue_6/0897.pdf

 

[DeeNos]

Dear researcher,

Thank you for sharing your findings and questions regarding the concentration of H3CO2 in a CO2+H2O reaction. Your research is very relevant, as understanding the chemistry of the oceans is crucial for understanding the effects of increasing atmospheric CO2 levels on our planet.

To address your first question about the equilibrium constant, you are correct that the equation [products]/[reactants] = K is used to calculate the equilibrium constant. In this case, the products are H2CO3 and the reactants are CO2 and H2O. However, it is important to note that the equilibrium constant is not a concentration, but rather a ratio of concentrations. This means that your calculation of 0.030 moles/L for the equilibrium in the forward direction is not the equilibrium constant, but rather the concentration of CO2 in the reaction.

To calculate the equilibrium constant, you will need to know the concentrations of all three species (CO2, H2O, and H2CO3) at equilibrium. This can be determined experimentally or by using an equilibrium expression and solving for the unknown concentration. For example, for the reaction CO2 + H2O = H2CO3, the equilibrium constant expression would be [H2CO3]/([CO2][H2O]). If you know the concentrations of CO2 and H2O at equilibrium, you can solve for the concentration of H2CO3 and then use that value to calculate the equilibrium constant.

Regarding your second question about the equilibrium of H2CO3 = H+ + HCO3- and HCO3- = H+ + CO3^2-, these are known as acid dissociation reactions and are related to the equilibrium constant for the reaction CO2 + H2O = H2CO3. The equilibrium constant for these reactions can be calculated using the same method as mentioned above.

In regards to your question about the solubility of CO2 in water, the value of 1.45 g/L at 25 degrees C and 100 kPa is for pure CO2 gas. This means that the pressure of CO2 alone is 100 kPa, and it does not take into account the rest of the atmospheric gases. However, in a real-world scenario, the partial pressure of CO2 in the atmosphere is around 400 ppm (parts per million), which is equivalent to 0.04 kPa. This is much lower than the sol