Carbon dioxide dissolves in water to form carbonic acid, which is primarily dissolved CO2. Dissolved CO2 satisfies the equilibrium equation.
CO2(g) <----> CO2(aq)
The acid dissociation constants listed in most standard reference texts for carbonic acid actually apply to dissolved CO2. K=0.032 M/atm
For a CO2 partial pressure of 4.4×10-4 atm in the atmosphere, what is the pH of water in equilibrium with the atmosphere?
Note: Since carbonic acid is primarily dissolved CO2, the concentration of H2CO3 can be taken as equal to that of dissolved CO2.
My online homework is due tonight and I am hopelessly lost on how to solve this problem. :(
Dr Bob is almost always right, don't be so salty ashley
Try this.
C = pK
C = 4.4E-4*0.032 = about 1.41E-5
H2CO3 ==> H^+ + HCO3^-
k1 = (H^+)(HCO3^-)/(H2CO3)
(H^+)= (HCO3^-) = x
(H2CO3) = 1.41E-5
Solve for x = (H^+) and convert to pH.
Are you plugging this into a data base? If so let me know how it turns out? And show your work when you do.
Yes I have online homework which I use to submit my homework on. I calculated H2CO3 and got the same value as yours. then I solved for x and got 2.3 x 10^-6 M. Finally i plugged it into the ph formula to get a ph of 5.64. Thank you so much! It was correct!
Hi,
Where did you get K to equal (H+)(HCO3-)/(H2CO3)?
General Ka expressions take the form Ka = [H+][A-] / [HA]. In this case A- is the anion HCO3- and HA is the acid H2CO3.
To solve this problem, we need to use the equilibrium constant expression for carbonic acid, which is equal to the ratio of products to reactants. Since the problem states that carbonic acid is primarily dissolved CO2, we can use the concentration of dissolved CO2 instead of carbonic acid.
The equilibrium expression for the dissolution of carbon dioxide in water is:
CO2(g) <----> CO2(aq)
The expression for the equilibrium constant (K) is given as 0.032 M/atm. This means that the concentration of dissolved CO2 divided by the pressure of CO2 in the atmosphere is equal to 0.032.
Now, we can start solving the problem.
1. Convert the given CO2 partial pressure to concentration:
The CO2 partial pressure in the atmosphere is given as 4.4×10-4 atm. Since concentration is in M, we need to convert atm to M. The ideal gas law equation can be used to convert the pressure to concentration:
PV = nRT
Rearranging the equation to solve for concentration (C):
C = n/V
Assuming 1 mole of CO2 occupies 22.4 L at STP (standard temperature and pressure), we can calculate the concentration:
C = (1 mol) / (22.4 L)
C ≈ 0.045 M (approximately)
Therefore, the concentration of dissolved CO2 is approximately 0.045 M.
2. Calculate the concentration of H2CO3:
As mentioned in the problem statement, we can assume that the concentration of H2CO3 is equal to the concentration of dissolved CO2, which is approximately 0.045 M.
3. Calculate the pH using the concentration of H2CO3:
To find the pH, we need to know the concentration of H3O+ ions, which can be calculated using the acid dissociation constant (Ka). The equation for Ka is:
Ka = [H3O+][HCO3-] / [H2CO3]
Since we are assuming the concentration of H2CO3 is equal to the concentration of dissolved CO2, we can substitute [H2CO3] with 0.045 M in the equation:
0.032 = [H3O+][HCO3-] / 0.045
Rearranging the equation to solve for [H3O+]:
[H3O+] = (0.032 * 0.045) / [HCO3-]
To simplify the calculation, we can assume that the concentration of HCO3- is approximately equal to the concentration of H2CO3 because carbonic acid is primarily present as dissolved CO2. Therefore:
[H3O+] = (0.032 * 0.045) / 0.045
[H3O+] = 0.032 M
Finally, we can calculate the pH using the formula:
pH = -log[H3O+]
pH = -log(0.032)
pH ≈ 1.5 (approximately)
Therefore, the pH of water in equilibrium with the atmosphere, given a CO2 partial pressure of 4.4×10-4 atm, is approximately 1.5.