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PH dependence of reduction of oxygen

  1. Jun 27, 2014 #1

    Qube

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    1. The problem statement, all variables and given/known data

    Why is this true?

    nc1vr.png

    Why is it that the reduction of oxygen is more favored in acidic solution rather than in basic solution?

    2. Relevant equations

    Acids are proton donors.

    3. The attempt at a solution

    Does this have anything to do with the fact that the reduction of oxygen is in part a protonation reaction? In other words, what I'm seeing is simply that protonation is more favored in acidic media than in basic media?
     
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  3. Jun 28, 2014 #2

    Qube

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    Anyone
     
  4. Jun 28, 2014 #3

    Borek

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    Exactly.

    Another (but equivalent) explanation will take Nernst equation into account. Just like kinetics depends on the concentration of species, potential is a function of concentrations as well.
     
  5. Jun 28, 2014 #4

    Qube

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    Sweet. Regarding the Nernst equation and how it predicts potential to decrease with increases in pH:

    Is the protonation of oxygen dependent on kinetic or thermodynamic factors? I'm guessing thermodynamic factors because in more basic medium, the thermodynamic potential for protonation decreases - i.e. the species in basic media are weaker acids than hydronium ion. In strongly basic solution the most viable acid is water, and that's not a particularly strong acid.

    Could kinetic factors also play a role?
     
  6. Jun 28, 2014 #5

    Borek

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    Nernst equation is about thermodynamic - it assumes equilibrium situation and fast reaction (which is more or less equivalent of doing the measurements with infinitely low current). When the reaction is slow observed potentials will vary from those predicted by the Nernst equation. That's where overpotentials come in - we have to increase the potential to speed up the reaction. That's why you will observe chlorine evolving from the solution of chlorides, even if potentials suggest oxygen should evolve first - oxygen is just pretty lazy and never reacts fast on electrodes :wink:
     
  7. Jun 29, 2014 #6

    Qube

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    Cool. So why does the potential decrease relative to the standard hydrogen electrode as the pH goes up for the oxidation of hydrogen protons? I'm guessing that it's easier taking hydrogen from hydronium ion in acidic solution than it is taking hydrogen from water in basic solution, since hydronium ion is less stable than water?
     
  8. Jun 30, 2014 #7

    Borek

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    They are not "easier to remove" - there are just more of them, so if anything, they are "easily available".
     
  9. Jun 30, 2014 #8

    Qube

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    Why not just take protons from water then? There are 50 plus moles of water even in most acidic solutions ... Excluding things like 18 M sulfuric acid.
     
  10. Jun 30, 2014 #9

    Borek

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    There is a difference between proton in a water molecule and proton in hydronium cation. The latter is much easier to remove.

    Note that because of water autodissociation hydronium is always present in the solution, just its concentration changes.
     
  11. Jun 30, 2014 #10

    Qube

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    Why much easier? Isn't it the same HO bond?
     
  12. Jun 30, 2014 #11

    Borek

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    No, its not the same. Besides, in the case of water you get an anion and a cation (so they attract) in the case of hydronium you get a cation and a neutral water molecule (so the attraction is between a charge and a dipole, and is orders of magnitude weaker).
     
  13. Jun 30, 2014 #12

    Qube

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    Gotcha. Does this have to do with the Ka values of hydronium ion and water?
     
  14. Jun 30, 2014 #13

    Borek

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    All these things are related in some way.
     
  15. Jun 30, 2014 #14

    Qube

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    So if I'm understanding you correctly the reduction of H+ in water solution to hydrogen gas is dependent on the availability of protons. It is easier to take a positive charge from a positively charged molecule rather than a positive charge from a neutral molecule, so the reduction of H+ in acidic solution is easier than the reduction of H+ in neutral solution, in which the main source of H+ will be water. And this is closely related to Ka values because Ka is a measure of the tendency of an acid to lose a proton to a base in water solution.
     
  16. Jul 1, 2014 #15

    Borek

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    Yes.
     
  17. Jul 1, 2014 #16

    Qube

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    Isn't the activity of the proton higher in acidic solution than in basic solution?
     
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