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Why do Manganate(VII) ions need scidic conditions?

  1. Jun 26, 2011 #1
    1. The problem statement, all variables and given/known data

    [itex]MnO_4^{-} + 8H^{+} + 5e^{-} \rightleftharpoons Mn^{2+} + 4H_2O[/itex]

    1. Explain why the presence of an acid is necessary for aqueous permanganate ions to function as an oxidizing agent.

    2. Give two reasons for the aqueous permanganate ions acting as an oxidizing agent in acidic solutions.



    3. The attempt at a solution

    The only thing I can think about is the similarity between this reaction and the Iodine-peroxodisulphate reaction where the repulsion between the negative ions prevents fast reaction unless positively charged [itex] Fe^{3+} [/itex] are present to catalyze the reaction.

    So my guess is that since permanganate ions are negatively charged and ions that need to be oxidized like [itex] Cl^{-} [/itex] are negatively charged they repel each other. So a positive ion - the [itex] H^{+} [/itex] ion is needed to make the reaction happen.

    I know this explanation is not perfect because permanganate ions also oxidize ions like [itex] Fe^{2+} [/itex].

    So any ideas?
     
    Last edited: Jun 26, 2011
  2. jcsd
  3. Jun 26, 2011 #2

    Borek

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    Staff: Mentor

    The question is invalid - permanganate acts as an oxidizing agent not only in acidic conditions, it also works in neutral and alkaline solutions, just products are different (compare http://www.titrations.info/permanganate-titration).

    Won't it be enough to take a look at the reaction equation and think in terms of LeChatelier's principle?

    --
     
  4. Jun 26, 2011 #3
    I got the question from an A-Level Chemistry textbook in a chapter on electrochemistry.

    The question may be referring to the specific reaction and why the presence of an acid improves the the oxidizing capabilities of permanganate ions.

    Does it act like a catalyst?
     
  5. Jun 26, 2011 #4
    I say this because the standard electrode potential for the above reaction is more positive than the standard electrode potentials for the reactions in neutral conditions.
     
  6. Jun 26, 2011 #5

    Borek

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    Staff: Mentor

    Do as I told you: take a look at the reaction equation, think in terms of LeChatelier's principle.

    What is catalyst definition? Is it consumed in the reaction?
     
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