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PH and Electrolysis

  1. Mar 10, 2006 #1
    Hey all. It has been a while.

    I have been interested in electrolysis for a while now and have been learning how different factors affect the reaction, e.g. distance of electrodes, concentration of solution, volume of solution electrolysed etc.

    However I am having problems finding research into how the pH should change during electrolysis (and not the affect the pH has on electrolysis). Because of this I have been researching it, practically, in chemistry (when the lab is free). I have been electrolysing copper (II) sulphate and testing the pH ever 5 minutes over a 45 minute period. However the results are not consistent and do not produce any sort of correlation.

    I was hoping someone could enlighten me to what should happen to the pH as any solution electrolyses.

    Cheers in advanced.

    Nice to see you all again.

    The Bob (2004 ©)
     
    Last edited: Mar 10, 2006
  2. jcsd
  3. Mar 10, 2006 #2

    GCT

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    while some reactions do produce a change in pH, others don't involve the acidity of the solution. The essence here is that, you've got to consider the equilibrium, if you're adding energy to produce the pH change, that is the energy affects the pH directly, than obviously you can model the situation to the experiment. If you are to produce any reactions which results in acid and base agents, so to speak, you'll have changes in pH. As for whether the direct electrical energy, somehow, alters the state of the autoprotolysis of water completely, upon reaching equilibrium, I'm not sure how that could work. If you were to change the temperature, the pH may change. Simply, a lot of factors affect the pH, but you should first consider the chemicals involved and then model the system.
     
  4. Mar 10, 2006 #3
    Thanks for that.

    I have a solution of copper (II) sulphate, so adding electrons (from electricity) would give:

    [tex]Cu^{2+}_{(aq)} + SO_4^{2-}_{(aq)} + 2e^- \rightarrow Cu_{(s)} + SO_4^{2-}_{(aq)}[/tex]

    But from this, I cannot see any reason for the pH to change. It would only change once the concentration of copper was very low and the hydrogen in the water was reducded to hydrogen gases. Then the pH would change but before that I can see no reason for it not to.

    Am I right or on the wrong track? [I ask mainly because I found the pH to change over time but the equation would not suggest it at all].

    Cheers

    The Bob (2004 ©)
     
  5. Mar 10, 2006 #4

    GCT

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    Copper can react as an acid, if you remember, although I'm not quite sure how efficient it'll be at this electronic state. How did the pH change with respect to time, did it increase/decrease, exponential, linear, no pattern....? Also, you've got to consider the activity pH, which should change, with the electrolyte composition.
     
  6. Mar 11, 2006 #5
    The pH decreased exponentially and seemed to level out. However I did stop the experiment before the hydrogen was produced at the anode.

    The Bob (2004 ©)
     
  7. Mar 11, 2006 #6

    GCT

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    You can't actually stop such reactions from occuring completely, so if the pH were to decrease, that is become more acidic, you would expect agents that are more or less acidic in the solution. The pH may have also changed slightly due to coprecipitation and perhaps the formation of hydroxide adducts, but none of these really affect the autoprotolysis equilibrium constant. You'll either need to identify an acidic agent, you can also account for the equilibrium change in pH in terms of activity pH. Really, I don't exactly know the details of this experiment, but you may also want to consider the existence of contaminants, silica glassware's affect on pH.

    you should also take the temperature into account. If you were taking the pH with the pH electrode while working simultaneously with an electrolysis electrode in close context, this may have also affected the functioning of the pH electrode.
     
  8. Mar 11, 2006 #7

    Borek

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    Copper reduction is only half of the processes taking part in the solution. What is the second half-reaction? Oxygen evolution? If so, here is answer to your question:

    2H2O -> 4H+ + O2 + 4e-
     
  9. Mar 12, 2006 #8
    Well ok. I meant I stopped the electrical current.

    Hum..... Problem is I don't know what could be. I was thinking of the water but I cannot see why the reaction would favour the water rather than the copper sulphate.

    That is a good point but I do not understand why it would change over time.

    This has been covered by using the Ph meter with a temperature setting.

    Well the second half should be the oxidation of sulphate ions. This should not have an affect on the pH. So the answer would seem to be the seperation of water. However, I do not see why water would be favoured over the copper sulphate (again). Is it to do with the electron densities of the two molecules/ions or simply the attractive charges? Or is it something else that is not as rubbish a suggestion as mine.

    Cheers.

    The Bob (2004 ©)
     
  10. Mar 12, 2006 #9

    Borek

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    No. Such a thing requires very high potential. Oxygen evolution goes at much lower potential, thus it is favoured as long as the water is present.
     
  11. Mar 12, 2006 #10

    GCT

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    Ok, my mistake, the copper ion should cause a basic solution, as possibly due to basic oxides, hydroxides, and hydrates. Thus removing copper, may have increased the acidity. I'm going not exactly certain about this though so I'm gonna research the exact details of the side reactions of copper in aqueous solutions.

    We may also need to consider the effects of charge transfer transition of copper in a water complex, and its affect on the basicity if any.
     
  12. Mar 12, 2006 #11

    Borek

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    It has nothing to do with the electrolysis. We start with the neutral solution of copper sulfate. Perhaps very sligthly basic due to hydrolysis, but this effect is very small, and even if it is present it changes starting pH only - and doesn't influence anything that happens later.

    What happens in this case is that on one electrode copper cation is removed from the solution while on the second electrode H+ are introduced, lowering pH. Once copper is removed you will be left with the solution of pure sulphuric acid, that's all.
     
    Last edited: Mar 12, 2006
  13. Mar 12, 2006 #12

    GCT

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    you've made a typo there, are you saying that the presence of sulfate will decrease the acidity?

    bob was trying to investigate the changes in pH, the changes in pH are influenced (whether to small or large degress) by the copper ions, you should investigate the original context of his query.
     
  14. Mar 12, 2006 #13

    Borek

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    Mea culpa, I meant solution prepared by the dissolution of copper sulfate in DI water can be slightly acidic due to copper ion complexation by OH- (which is in effect identical to weak base hydrolysis).

    Please reread his original post. He asked about the changes of pH during electrolysis of copper sulfate:

    and the answer is pH goes down, as the copper ions get replaced by H+.
     
  15. Mar 12, 2006 #14
    Well... thank you both for your replies. Though I had my own theories (as you have seen and corrected), your input has really helped my understanding. Allow me to ensure I understand.

    The solution should start neutral, ignoring the affects of hydrolysis. Once electrolysis starts, copper is obvisouly produced at the cathode. However oxygen is formed at the anode because it has a lower electrode (?) potential than sulphate ions. The reaction that forms the oxygen (2H2O -> 4H+ + O2 + 4e-), as Borek stated, means that the concentration of hydrogen ions increases and, thus, increases the pH.

    Now this is all good, we are left with sulphuric acid and copper on the cathode. But does the copper actually make any difference to the pH itself? I mean does the liberation of copper actually change the pH, assuming no oxygen was formed? I beleive it wouldn't but that is only based on common sense (and as I have learned, common sense is nothing in chemistry).

    Cheers for all your help so far.

    The Bob (2004 ©)
     
  16. Mar 12, 2006 #15
    Sorry for the spelling mistakes in post 14 but my computer will not allow me to edit my post.

    The Bob (2004 ©)
     
  17. Mar 12, 2006 #16

    Borek

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    No. If you can substitute some other electrode reaction pH will not change. For example if you will use zinc anode - zinc will dissolve in the solution (effectively replacing copper) and the pH will not change.
     
  18. Mar 13, 2006 #17

    GCT

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    Bob, now that you have your probable causes, you'll need to establish them using some elaborate experiments. There's no sense in doing the experiments, and then asking the questions here to completely resolve your queries. It's most often the case that bringing up a subject as vague as this one will produce different opinions.

    Try to specify the question that you have, so that you can work experimentally to arrive at a definite conclusion.
     
  19. Mar 17, 2006 #18
    The zinc will replace the copper? From an electrode? Now I am lost. So the copper will not change the pH because the copper electrode will replace the copper? So why does the mass of the cathode change?

    Thanks for that. Once I understand my last question then I will redo some tests and see what happens.

    Thanks for all of your help. The both of you.

    The Bob (2004 ©)
     
  20. Mar 17, 2006 #19

    Borek

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    Look at the situation this way. Solution must be neutral - same amount of cations and anions. If you remove Cu2+ some other cation must enter solution. It can be two H+, or Cu2+ dissolved on the anode, or Zn2+ dissoved on the anode - it all depends on what material anode is made off. If the anode dissolves, pH doesn't change. If anode doesn't dissolve, oxygen evolves and H+ enters the scene, changing pH of the solution.
     
  21. Mar 17, 2006 #20
    Yer, I see what you mean.

    Cheers Borek.

    More than likely this thread will be approached from time to time. :smile:

    Cheers all. :biggrin:

    The Bob (2004 ©)
     
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