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Why does saltwater increase rate of corrosion?

  1. Jan 16, 2017 #1
    Why does salt water increase rate of corrosion, really? Most common answer I read was that salt makes water a better electrolyte. No further explanation.

    However, I'm not really getting the mechanism of this. I understand the mechanism of pitting corrosion in presence of chloride ions and I understand how chlorides break down the passive layer. Also, I'm ok with metal chlorides being more soluble so they do not form non-soluble layer of protection for protection from further corrosion.

    However, how does being a better electrolyte make up for higher corrosion rate? In regular NaCl solution, for regular DC current, conductivity is achieved by chloride ions and hydrogen ions attraction towards charged electrodes and redox reactions (and transferring electrons through the loop), not by moving charges themselves. So how does conductivity make up for faster rusting, is there a reaction, a mechanism, an electron transfer? Where?
     
  2. jcsd
  3. Jan 16, 2017 #2

    Borek

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    Lower resistance of the solution definitely helps to speed up charge transfer, especially when redox centers (oxidation/reduction) are separated (that is, reactions don't take place in exactly the same spot).
     
  4. Jan 17, 2017 #3
    But where does electrolyte conductivity take place here? If I have a reaction, for example, Fe ---> Fe2+ + 2e- , electrons move through the metal itself to the place where they will get "spent": H2O + (1/2)O2 + 2e- ---> 2OH- .

    One other thing, as I see it, free Na+ and Cl- ions help to disrupt H2O molecules in a sense, encouraging them to also split into their respective H+ and OH- ions more often. Since there is higher concentration of OH- ions, OH- is getting removed from the solution by precipitating Fe2+ + 2OH- ---> Fe(OH)2, therefore shifting cathodic reaction to the right, and speeding up the process?

    Second thing that comes to my mind is that chloride ions from the solution "remove" iron cations by reaction Fe2+ + 2Cl- ---> FeCl2 and therefore, shift reaction Fe ---> Fe2+ + 2e- to the right, furthermore speeding up the oxidation process. Is this correct?
     
  5. Jan 17, 2017 #4

    Borek

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    This circuit is not closed yet.
     
  6. Jan 17, 2017 #5

    Borek

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    In general - they don't. Sure they use water molecules for solvation, but the effect on the water autoionization is negligible.

    To some extent complexation helps, no doubt about it. But that's less important. We are more interested in the corrosion kinetics than in the final equilibrium.
     
  7. Jan 17, 2017 #6
    I think this is the part that's bugging me. Why does the circuit even need to be closed? Fe2+ keeps getting disolved and OH- keeps getting in the "solution". Do Na+ and Cl- work like a salt-bridge in a galvanic cell in a way, enabling to somewhat neutralize localized (anodic) positive charge (from Fe2+), and (cathodic) negative charge (from OH-) distribution ?
     
  8. Jan 17, 2017 #7

    Borek

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    If the circuit is not closed you have a buildup of charge.

    Let's start with a simplified model. We know both half reactions occur in separate places (not necessarily far from each other, but separated). Before the reaction:

    Fe -------------------------- O2

    System is neutral on both ends. Now the reaction proceeds and electrons move. After the reaction:

    Fe2+ -------------------------- O22-

    There are charges on both ends, and they will prevent further reaction, as electrons now have to move against the electric field (you can also think in terms of a potential difference being created in the system) which is much harder (requires much more work done). That's where closing the circuit comes in handy - it helps to discharge the system so that both ends become neutral again and the reaction can proceed further.
     
  9. Jan 17, 2017 #8
    Great, thanks!
     
  10. Jan 17, 2017 #9

    Nidum

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    Do you mean flowing sea water or static lab salt solutions ?
     
  11. Jan 17, 2017 #10
    Static.
     
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