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Why cations migrate toward cathode?

  1. Nov 28, 2015 #1
    2drcv1i.png

    Why Na+ (which is positive in charge) migrate to the cathode which is also positive in charge??
    They should repel each other right?
     
  2. jcsd
  3. Nov 29, 2015 #2

    symbolipoint

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    Cupric cations are becoming reduced at the cathode, and the sodium cations, being mobile in the gel and then in solution, are replacing the cupric ions as they become reduced and deposit.

    THE CATHODE SUPPLIES NEGATIVE CHARGE. IT IS NOT POSITIVE.
     
  4. Nov 30, 2015 #3
    https://en.wikipedia.org/wiki/Anode

    Anode sorta means current flows away, but originally when charge was explored, it was presumed that the positive particles moved (only later was it discovered that in an electrical circuit, it was the negatively charged electron that moved). It's also a source of confusion in electrical schematics where arrows generally point in the direction opposite of the electron flow.
     
  5. Dec 1, 2015 #4

    DrDu

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    Of course Cathode is positive in the setting discussed above.
    However, the potential of the cathode is not relevant for the motion of the sodium ions in the bridge.
    There is a double layer at the cathode solution boundary, i.e. there are positive charges at the surface of the cathode and negative counter ions at the solutions boundary. Hence the solution is at a lower potential than the cathode.
    With the same line of argument, the solution on the anode side is at higher potential than the anode.
    When there is no flow of current, the two solutions are at equal potential. However, if the battery is short cut, then the two electrodes will be at same potential and there will be a potential drop from the anode to the cathode which drives the ionic current in the salt bridge.
    The real situation with a resistor between the two electrodes will be somewhere in between, i.e. part of the voltage will drop on the resistor and part of if on the salt bridge.
     
  6. Dec 5, 2015 #5
    Cathode is negatively charged. It's called cathode because cations migrate there.
     
  7. Dec 5, 2015 #6

    DrDu

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    The charge of the cathode is different in electrolytic cells and galvanic cells. The cathode is defined to be the electrode where reduction takes place.
     
  8. Dec 5, 2015 #7

    Merlin3189

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    As a non-expert bystander can I add a couple of observations?

    The voltmeter shows the Copper cathode is positive wrt the Zinc anode, whatever you call them and whatever theory says.

    Do we know that the Na+ and NO3- migrate? I would have thought that current in the salt bridge would more likely be carried largely by the more mobile H+ ions and OH- ions. I would expect that any net movement of Cu, Zn, NO3 ions would be diffusion along a concentration gradient, possibly biased a little bit by the very small electric field. With no concentration gradient, would Na+ migrate at all?
     
  9. Dec 5, 2015 #8

    DrDu

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    Usually, the salt bridge solution is more or less neutral. Hence the concentrations of H+ or OH- is orders of magnitude below the concentrations of Na+ and NO3-, hence, although the mobility of H+ is higher than that of the other ions, the contribution to the current is negligible.
    Also, the electric field isn't very small. Let ##R_i## be the internal resistivity of the salt bridge, ##R_e## the external resitstivity i.e. the load of the circuit, and ##U_0## be the voltage of the galvanic cell, then the voltage drop over the salt bridge is ##U=U_0 \frac{R_i}{R_i+R_e}##.
     
  10. Dec 5, 2015 #9

    Merlin3189

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    Thanks for the comment DrDu. I take your point about the low concentration of H+ and OH-: I shall have to do some sums on this. As for the field I was suggesting that U0 of 1.2V means Usalt bridge cannot be greater and the field might only be about 10 V/m, which I thought small. Again, I need to do some sums before I say anything more.

    Looking for some other discussion of this point, I came across,

    The Royal Society of Chemistry has a http://media.rsc.org/Classic%20Chem%20Demos/CCD-43.pdf [Broken] which, if it works as described, seems to suggest rapid ion migration due to an electric field. I'm surprised that they don't provide a video, or photographs, so shall have to try this one myself. But it is based on the movement of OH- ions and I think H and OH may be special cases. (Because H and OH entities are ubiquitous in aqueous solutions, might an H+ or OH- move virtually by transfer of charge and be therefore hyper-mobile compared to other ions?)

    And re. my previous comment, a quote from "All About Electrochemistry" in "Chem 1 Virtual Textbook"
    But this is not an area I'm familiar with, so I must leave this as my speculative comments until I have done some more investigation.
     
    Last edited by a moderator: May 7, 2017
  11. Dec 6, 2015 #10

    DrDu

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    There are tables of ionic mobilities e.g. on the German page about ionic mobility: https://de.wikipedia.org/wiki/Ionenbeweglichkeit
    So the current in a salt bridge of area A and length L (i.e. volume V=AL) is ##I=A\mu c F U/L## where c is the concentration and ##F\approx 100000 C/mol## Faradays constant. Assuming the salt bridge to consist of a cube of 1 cm^3 of a 1 molar solution, we get for the resistance ##R=1/(A (\mu_{Na^+}+\mu_{NO_3^-})c F/L)=1/(10^{-4}(13.6\times 10^{-8}) 10^{3} 10^5/10^{-2})## Ohm which would be about 7 ##\Omega##.
     
  12. Dec 6, 2015 #11

    Borek

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    Yes, mobilities (actually we measure something called "limiting ion conductivity") of H+ and OH- are much higher than mobilities of other ions. Not orders of magnitude, but several times higher.
     
  13. Dec 6, 2015 #12

    DrDu

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    about 7 and 4 times larger, respectively than that of other singly charged ions, on the mean.
     
  14. Dec 6, 2015 #13

    James Pelezo

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    Can you share a reference that supports your claim as it applies to Galvanic Process?
     
  15. Dec 6, 2015 #14

    James Pelezo

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    I'll probably be castigated for this, but there seems to be a considerable misunderstanding in the function and chemistry of electrochemical cells, when they are actually very simple... Please forgive my being wordy, but in the spirit of supporting the forum an chemical education, I would like to add my interpretation of basic concepts about electrolysis and galvanic processes for those who would wish to know how such systems function. It does not presume to include all details of explaining electrochemical processes; just basic concepts.

    For both Electrolytic and Galvanic/Voltaic Cells the electrodes in electrochemical cells are defined in terms of the chemistry that takes place at a specified electrode; i.e., 'Reduction Rxn' = Cathode and 'Oxidation Rxn' = Anode.

    Galvanic Cells:
    For instructional purposes, there are two Galvanic processes; 'Controlled Galvanic Process' and 'Uncontrolled Galvanic Process'. The Uncontrolled Galvanic Process is one in which oxidation and reduction reactions occur simultaneously in one cell. Example: Given a container of Copper(II) Sulfate ions in solution and inserting a Zinc metal bar directly into the solution results in reduction of Cu+2(aq) ions to Cuo(s) sticking to the surface of the zinc bar. Over time, the copper sticking to the zinc bar will form a coating that will prevent further reduction of copper ions to copper metal and the Galvanic/Voltaic process stops. For a controlled Galvanic/Voltaic process, the chemical processes are separated into anodic and cathodic cells allowing the system to sustain a continuous current of electric charge flow until the anodic material completely dissolves into solution leading to a 'dead battery'.

    Example of Controlled Zn/Cu Galvanic/Voltaic process (Refer to the diagram at the beginning of this thread):
    The Galvanic/Voltaic Cell for the copper/zinc system is a spontaneous reaction process giving current flow when connected. Copper ions in solution are reduced at the copper bar electrode by gaining (reduction) 2 electrons... Cu+2 + 2e- => Cuo(s) leaving the copper electrode deficient in electrons => positive electrode (cathode). In the zinc cell side, The Zinc electrode is being oxidized to Zinc(II) ions that are delivered into solution. Zno(s) => Zn+2 + 2e-. This oxidation half reaction leaves excess electrons in the Zinc bar and => negative electrode (anode). The voltaic cell will discharge until all of the anode is dissolved and no oxidation half reactions occur and the cell is a 'dead battery'.

    As for the salt bridge, its function is to maintain balance of charge as the Galvanic process proceeds. In the anode side of the cell where oxidation is occurring, there is an increase in positive charge due to the cations being delivered into solution. The Negative ions of the salt bridge therefore migrate to the anodic cell to neutralize the build up of positive charge. The Positive ions of the salt bridge migrate to the cathode cell to replace the positive charge loss when cations in solution are reduced to neutral causing a loss of positive charge in the cathodic cell solution.

    Electrolytic Cells:
    The electrodes in the Electrolytic Cell are defined in the same way; i.e., Oxidation => Anode & Reduction => Cathode. The difference is the chemistry of the ions in the solution cause the anode to assume a positive charge and the cathode to assume a negative charge; opposite that of the Galvanic Process. Example: NaCl(melt) => Na+(l) + Cl-(l). The Electrolytic Cell is non-spontaneous and requires an outside potential to drive the reactions and is therefore connected to a Galvanic Cell (battery) giving one electrode in the electrolytic cell a positive charge and the other a negative charge. The Na+ ions migrate to the electrode connected to the (-) electrode of the battery and undergo reduction by gaining an electron (Na+ + e- => Nao(s) and the chloride ions undergo oxidation to give chlorine gas (Cl2) ... 2Cl- + 2e- => Cl2(g). Commercially this is referred to as a Downs Cell.
     
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