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Topic: Titration (given pH of unknown acid?)  (Read 2734 times)

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Offline naffity

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Titration (given pH of unknown acid?)
« on: March 31, 2010, 09:02:05 PM »
"A student places 25.0 mL of an acid solution into an Erlenmeyer flask. The acid solution has an initial pH of 1.60. The student then titrates the acid to the equivalence point using a solution of NaOH. If 15.8 mL of base were required for complete neutralization, calculate the concentration of the base."

I thought the formula was Macid * Vacid = Mbase * Vbase so how do I find the concentration of the acid? I'm assuming you're supposed to use the pH, but how?

And on a semi-related note, when I do a titration lab, when do I know the solution has titrated? Today we did a "practice lab" with hydrochloric acid and sodium hydroxide with phenolphthalein as an indicator. It's supposed to turn really light light pink when it's done...but even if you add too much NaOH, it still eventually fades from dark pink to a light pink. If you add to little NaOH, there's a flash of pink that disappears really quickly. So when exactly are you done? Sorry if my incoherent rambling doesn't make sense; if you want me to clarify more later I can.

Umm yeah, that's about it. Please answer! I have an unhappy grade in chem! :[

Offline AWK

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Re: Titration (given pH of unknown acid?)
« Reply #1 on: April 01, 2010, 02:29:40 AM »
You need a concentration of base.

Offline Borek

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