1. 0.135M in HBr and 0.135M in HCHO2
2. 0.130M in HNO2 and 8.0×10−2M in HNO3
3. 0.180M in HCHO2 and 0.22M in HC2H3O2
4. 4.5×10−2M in acetic acid and 4.5×10−2M in hydrocyanic acid
Could you also please explain the reasoning behind the calculations (ie what’s a strong/weak acid/base)? I feel a bit silly and rather stumped. Thank you in advance!
The answers:
1. .870
2. 1.10
3. 2.24
4. 3.05
Good luck everyone! Mastering Chem cannot conquer the people of Yahoo! Answers.
2 Answers
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In each case, you act as if the solution contains only the stronger of the two acids.
1) The solution pH would be calculated based on the molarity of the HBr only. The formic acid would be ignored.
2) Use the pH of the NHO3; ignore the HNO2
3) This one is interesting since the Ka of formic acid is on the order of 10^-4 and acetic acid is on the order of 10^-5. Usually, the acids are not as close as this. However, to a first approximation, I believe that you would still go with the stronger acid, which is the formic acid. However, I could be wrong. I’d like to see this one done by someone more expert than myself.
4) The acetic acid is by far the strong. HCN has a Ka in the region of 10^-10, making acetic acid approx 100,000 times a stronger acid.
The idea is the the stronger acid releases more H3O^+ into the solution and by so doing suppresses the contribution of H3O^+ by the weaker acid. This is an example of LeChatelier’s Principle.
By the way, the four calculations above fall into two categories. Calculate 1 and 2 as a strong acid, where you would take the -log of the concentration of the acid. In 3 and 4, you must perform a weak acid calculation using the Ka of the acid. In 3, the formic acid is stronger than the acetic, but it still remains a weak acid, only partially ionized in solution.
