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KMnO4 Titration, why does it need to be in an acid?

  1. Oct 13, 2013 #1
    1. The problem statement, all variables and given/known data

    Fe(NH4)2(SO4)2*6H2O is mixed with H2SO4 and then titrated with KMnO4 until the equivalence point is reached.
    The question I am confused with is:
    What might have been the product(s) in the original solution if it had remained neutral? (if the solution was not acidified with H2SO4) How could you determine this?


    2. Relevant equations

    Net Ionic Equation:
    8H+ + MnO4- + 5Fe2+ => Mn2+ + 4H2O + 5Fe3+


    3. The attempt at a solution

    From the wikipedia page on this, it says that potassium permanganate degrades into MnO2 when reacted in a neutral equation. I would think that the redox reaction would still occur between the Fe2+ solution and the KMnO4, but it wouldn't reach the equivalence point since there is no acid...so it wouldn't change color?

     
    Last edited: Oct 13, 2013
  2. jcsd
  3. Oct 13, 2013 #2

    symbolipoint

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    Homework Helper
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    Gold Member

    You have written part of the justification for doing the titration in acid conditions. Hydronium ions participate in the reaction.
     
  4. Oct 13, 2013 #3
    Okay, so the above reaction couldn't take place without the hydrogen ions. Would a different reaction occur? Or would no reaction occur? My best guess would be something like: Fe2+ + MnO4- => MnO2 + O2 + Fe3+?
     
  5. Oct 14, 2013 #4

    Borek

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    Staff: Mentor

    Depending on conditions permanganate gets reduced to manganate, manganese dioxide, or Mn2+. Low pH guarantees only one reduction product and high enough redox potential to proceed with oxidation of Fe(II).
     
  6. Oct 15, 2013 #5
    I think the potassium manganate (VII) must be acidified, so that the hydrogen ions will mop up any excess oxygen atoms that were otherwise not taken up.
     
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