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Topic: Titration Problem  (Read 3005 times)

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Offline dantheman00114

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Titration Problem
« on: April 15, 2009, 05:06:55 PM »
In order to standardize a KMnO4 solution, 0.3301 g Na2C2O4 was dissolved in 30 mL water and 15 mL 3.0 M H2SO4. The KMnO4 solution was added to the Na2C2O4 solution until a pale pink color persisted. The titration took 31.1 mL of KMnO4 solution. What is the concentration of the KMnO4 solution?

I know that I'm probably going to have to start with a net ionic equation for the first addition of Na2C2O4 and H2SO4 with the sulfuric acid acting as the polyprotic acid.  Would I first have to convert to moles, and then figure out the concentration of the ions formed that will react with KMnO4 in the next reaction?? But then how do I write the equation and solve the equilibrium after the KMnO4 has been added?

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Re: Titration Problem
« Reply #1 on: April 15, 2009, 05:36:59 PM »
The only reaction you should worry about is KMnO4 + Na2C2O4.
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Offline dantheman00114

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Re: Titration Problem
« Reply #2 on: April 15, 2009, 06:50:00 PM »
why does the H2SO4 not affect the problem? should i be dealing with the total volume of 45 mL then?
and in that regard, wouldn't the problem just simply be a M x V = M x V problem??

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Re: Titration Problem
« Reply #3 on: April 16, 2009, 02:52:37 AM »
Acid is there only to lower pH so that reaction between permanganate and oxalate proceeds fast.

This is a potentiometric (redox) titration.
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