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Thread: Hydrogen Peroxide Decomposition

  1. #1 Hydrogen Peroxide Decomposition 
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    I am trying to write a lab report and I stupidly picked a very difficult topic which we have not covered at all. After lots of research, I am trying to write the simplest, acceptable hypothesis that I can manage:

    Basically, I need to explain why transition metals are better catalysts than group 2 metals.

    I've got it basically done, except someone told me:

    Group 2 metals do not make good catalysts for two reasons:
    1. It has only one accessible oxidation state and lots of catalytic mechanisms involve oxidative addition or reductive elimination
    2. Their d orbitals do not interact with organic ligands well. Since most metal mechanisms involve coordination chemistry this prevents group 2 from being able to do much.

    I understand accessible oxidation states after A LOT of reading.

    Could someone please tell me whether the decomposition of hydrogen peroxide involves oxidative addition/reductive elimination or whether it is number 2: coordination chemistry and all that complicated d orbital stuff which I would have to go research if it does?


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  3. #2  
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    Off the top of my head, I don't know exactly how metals catalyze the decomposition of H2O2.

    But, I'm afraid you might be a little confused about some things, because you mention oxidative addition and reductive eliminations, and then refer to "that d orbital stuff" as if it was something different. d orbitals are the main orbitals involved in most interactions between metals and organics in oxidative addition and reductive elimination. The two subjects are very closely linked.


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  4. #3  
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    Okay thank you very much!

    That was extremely useful, I guess it's more reading for me..

    :-D :-D
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  5. #4  
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    Hmmm, where exactly does hydrogen peroxide come from in all this? Knowing that might help. Hydrogen peroxide decomposes spontaneously to water. Oxidative addition and reductive elimination, as far as I know, involve the loss or gain of twosubstituents and a change of charge by 2. I suppose h2o2 losing an oxygen would increase the charge by two but I don't know that it would be reductive elimination. Hydrogen peroxide is a common byproduct in catalysed reactions, could you perhaps be getting confused. Not sure that this helped you with your problem at all unfortunately, but I can give you a few pointers when you are thinking about this sort of thing.

    Firstly you should look at the definition of a catalyst. A catalyst by definition is unchanged at the end of a reaction.

    Next you should look at the fact that during an oxidation reaction, when something is oxidised something else must be reduced to make up for it. So if you have a +2 metal centre in your catalyst, it will drop down to say 0 when you oxidise the substrate. (The metal cntre then needs to be put back to 0 somehow if the reaction is catalytic)

    As you can see already, two seperate metal oxidation states were needed for this reaction. Since 2nd row elements only have one stable oxidation state, then they cant possibly be suitable as catalysts.

    You might want to go on from there to look at the reasons why transition metals have more possible oxidation states, but I couldn't be bothered saying much more, and I suspect that complex understandings of the energy levels of different orbitals might be boynd the scope of your course.
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