[webguy54]

Q: Calculate the pH of a solution prepared by diluting 3.0mL of 2.5M HCl to a final volume of 100mL with H2O.

A: pH = 1.1

I first found the pH of 2.5M HCl by using: pH = -log [H+].  I calculated a pH of -0.398.  I then substituted this into the Henderson-Hasselbalch equation as the pKa: pH = pKa + log ([base]/[acid]).  For the ratio within the log, I used 97% for the base (water) and 3.0% for the acid (HCl).  I used these numbers because the ratio was out of 100mL = 3.0mL(acid) + 97ml(base). After calculating I got the correct answer of pH = 1.1

Did I do this problem correctly or did I simply get lucky?  I understand the steps that I performed...I didn't just plug in numbers, but I feel as if I could have done it differently.  Is there a different way to calculate this problem?

Thank you

[Borek]

did I simply get lucky?

Yes, you are completely off I am afraid.

Henderson-Hasselbalch equation is for weak acids, you have solution of strong acid, so it is 100% dissociated. Calculate concentration of the acid after dilution - and as the acid is fully dissociated, concentration of H⁺ will be that of acid itself...