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Why do we use electrode potential to predict other reactions

  1. Sep 11, 2016 #1
    Probably the title wasn't really clear. What I am asking for is an explanation.
    Isn't the potential of the electrode is created by a combination of ions and electrons? and we can't possibly measure the absolute potential. However, we can measure it relative to other electrodes such as hydrogen.

    So why do we use the electrode potential to predict if a reaction will happen or no? In this example, imagine you have a bowl made out of ##Fe## and you have ##Cu^{+2} ## in it. Obviously a reaction will happen.. and you can predict that using the standard potentials. But as I said earlier, The potential is created by a combinations of ions and electrons. However we have only ions here.

    So does the electrode potential actually refer to how much voltage an electron gains when it moves from infinitely far away to an ion? and the opposite happens if it is reduction?

    Why is that? I also know that voltage is an intensive property for electrode.

    P.S The problem with this topic is not that I am not able to solve its problem but me not knowing what happens in atomic scales
     
  2. jcsd
  3. Sep 11, 2016 #2

    Borek

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    Staff: Mentor

    You don't need electrons for the potential difference to exist, existence of any charge is perfectly enough.
     
  4. Sep 11, 2016 #3
    Yea that is fine of course.

    But the electrode potential is exactly equal to the potential if you only had positive ions? ( Or even guess a single ion maybe?)
    Also, Please help me with this. When does the galvanic cell stop? These question and their answers irritates me.
    Some websites dont mention this information, Others say it will stop when the electrodes are depleted and finally some says when the equilibrium is reached.
    and I have checked like 30 websites before asking this :/
     
  5. Sep 11, 2016 #4

    Borek

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    Staff: Mentor

    Doesn't matter what kinds of charge are present, the potential is always the same thing.

    To some extent both answers mean the same. For redox reactions equilibrium is typically shifted so far in one direction that for all practical purposes we can assume one of the electrodes (or more generally one of the reactants) was depleted.

    But you will get much better explanation in thermodynamics - every reaction (electrochemical reaction included) goes till the Gibbs energy of the system reaches its minimum possible value. Then the system is at equilibrium and progress of the reaction stops.

    I have a feeling part of your problems may stem from the fact you don't know rather basic things about chemistry and physics, so you feel lost and you don't see obvious connections when presented with the explanations. You will probably do yourself a favor starting from the beginning, GenChem101, any general chemistry book.
     
  6. Sep 11, 2016 #5
    I wouldn't say it stems from physics. The problem is we haven't took equilibrium nor gibbs free energy. Some of these topic we will take later this semester, other are in the 2nd semester.

    I am pretty sure I know my physics and chemistry pretty well. I have never spent on a topic this much time trying to understand it atomically. Mostly I just read the book and go solve all of the old exams problems and I did that with this one. It was all correct.

    But I just cant get my head around this in atomic scales.

    I just want a detailed physical view just as you would in circuits using Drude model which "Matter and interactions book" deeply used to explain how the fields are in the wire, how the surface electrons sustains the electric field ...etc That Is what I was searching for in the internet for the past 4 days but I couldn't find any.

    A good start would be if you imagine two electrodes zinc and copper. Zinc loses electrons more readily so you have a bigger negative charge build up on the zinc electrode than copper and zinc ions are near the surface of the electrode them together produce an electric field, From there, I need someone to walk me through how an electrons travel through the galvanic cell and how fields affect of both electrodes affect them.

    All these problems rises to me because I like to think of the contribution of every single electrons and how it affects the overall potential. If you could give me a detailed physical view instead of a chemical point of view, I am pretty sure that would solve all

    P.S I was asking for if you have lets say x electrons and you move an electron from infinitely far away to these electrons, the voltage will be higher than if you move that electron to only 1 electron. This idea doesn't seem to apply to electrodes. That is why I need a clarification

    And thanks for answering the previous question :D. Can you please help me with this?
     
    Last edited: Sep 11, 2016
  7. Sep 12, 2016 #6

    James Pelezo

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    Gold Member

    I might suggest looking at the topic in terms of the ‘Activity Series’ as applied to Single Replacement Reactions; i.e., A + BX => AX + B. Element A will replace element/ion B because A undergoes oxidation (loses electrons) more easily than B; or, B is more resistant to oxidation and prefers to exist in a reduced state in the presence of A.

    Example…

    Given the elements in order from most active (ease of oxidation=reducing agent) to least active (resistant to oxidation=oxidizing agent):

    Li > Rb > K > Mg > Zn > Cr > Fe > Ni > Cu > Ag > Au

    <= Stronger Reducing Agent upload_2016-9-12_19-42-49.png Stronger Oxidizing Agents =>

    The potential for ‘spontaneity’ is the stronger reducing agent will replace the oxidizing agent. The oxidizing agent (Cu+2) prefers to undergo reduction in the presence of Zno and converts to its basic elemental standard state, Cuo(s) … Zno(s) + Cu+2(aq) => Zn+2(aq) + Cuo(s).

    You might also compare the valence shell electron configurations of one metal against another to theorize why one metal is more readily undergoes oxidation while the other prefers the reduced state in the presence of the reducing agent.


    Zn[Ar]3d104s2 => Zn([Ar]3d104s0)+2 + 2e- (Oxidation)

    Cu[Ar]3d94s2 <=> Cu[Ar]3d104s1 => Cu([Ar]3d104s0)+1 => (Cu([Ar]3d94s0)+2 + 2e- => Cu[Ar]3d94s2 <=> Cu[Ar]3d104s1 (Reduction)

    ( The incomplete 4s orbital of Cu+2 would be paramagnetic in the presence of Zinc and attract electrons from Zinc to form a more stable orbital configuration => ‘Reduced form of Copper in Basis Std State) … Theoretically.
     
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