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thinking about oxidation and reduction, what happens to the silver metal and the silver ions?

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    $\begingroup$ How would you tell the difference between "yes, it would occur" and "no, it wouldn't"? $\endgroup$ – Ivan Neretin Apr 3 '20 at 4:38
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    $\begingroup$ @IvanNeretin use Ag-110m; compare diffusion of gold in gold $\endgroup$ – user7951 Apr 3 '20 at 7:29
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For OP: As I have realized from James Gaidis' answer, you were very minimalistic in describing context of your question, what is very bad practice, as much more burden is put on shoulders of responders via the question interpretation and clarification. As I have just now realized, you may, or may not speak about a solution.


Yes, the reaction will occur, and no, the reaction will not occur.

There is ongoing equilibrium reaction $\ce{Ag <=> Ag+ + e-}$ with zero net result.

Initially, when metallic silver is inserted, very tiny net result of the above reaction happens, until there is achieved the equilibrium potential. It is determined by the initial potential before inserting into the solution and by the effective capacitance of the electrode.

If you charged positively ( e.g. electrostatically ) a piece of metallic silver before inserting into the solution, a very small amount of silver would dissolve. If negatively, a very small amount would deposit itself.

As @MaxW has mentioned in his answer, isotope tracing is frequently used to track the reaction mechanisms of complex processes and also kinetics of systems in equilibrium, that seem macroscopically static.

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    $\begingroup$ This is a very important and subtle point you made. I wish all teachers like electrochemistry like your teachers. East Europeans were too good in electrochemistry. $\endgroup$ – M. Farooq Apr 3 '20 at 4:54
  • $\begingroup$ :-) Hehe. But the truth is, I liked electrochemistry at university. $\endgroup$ – Poutnik Apr 3 '20 at 4:55
  • $\begingroup$ Oops... Took to long to type my answer. $\endgroup$ – mpprogram6771 Apr 3 '20 at 5:03
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    $\begingroup$ @mpprogram6771 That is OK. The quick answers are offen overcome by later thorough answers. :-) $\endgroup$ – Poutnik Apr 3 '20 at 5:05
  • $\begingroup$ @Poutnik, You have made really important point, "If you charged positively ( e.g. electrostatically ) a piece of metallic silver before inserting into the solution, a very small amount of silver would dissolve. If negatively, a very small amount would deposit itself." Have you seen this described in your textbooks. $\endgroup$ – M. Farooq Apr 3 '20 at 14:50
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Between the solution of silver ions and the silver atoms on the electrode there is ongoing exchange of silver ions for silver atoms with a zero net change. Since Ag mined from the earth has no radioactive isotopes, it would possible to detect the change by using a radioisotope of silver say $\ce{^{108\mathrm{m}}Ag}$ in the solution. After immersing the Ag electrode in the radioactive solution it too would become radioactive indicating that a exchange did happen.

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Short answer: No

Long answer: Honestly, there's really no good way to tell.

Think about it this way: the only plausible redox reaction between the Silver(I) cations and the Silver metal is the exchange of electrons on the surface of the metal. And if this happened, the $\ce{Ag(s)}$ would donate one of it's valence electrons to $\ce{Ag+(aq)}$, resulting in a $\ce{Ag+(aq)}$ ion and a $\ce{Ag(s)}$ atom. See the change? there isn't one. So as far as we know, the surface of that silver electrode could constantly be exchanging electrons with the solution of ions, but we would never know, and It really has no significance whether it does or not, because it wouldn't affect much.

Credit to @Ivan Neretin for prompting this answer.

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I'm going to be a literalist, and then an extremist, and then a speculator.

If you put a silver wire into silver nitrate (which is literally a solid, mp = 210 C), no visible reaction will occur.

However, silver nitrate decomposes when heated: 2 AgNO3(l) → 2 Ag(s) + O2(g) + 2 NO2(g) Qualitatively, decomposition is negligible below the 210 C melting point, but becomes appreciable around 250 °C and it totally decomposes at 440 °C. (Wikipedia)

So, if you warm the silver nitrate solid until it is extremely melted, it may be enough of an oxidizer to oxidize silver metal to the oxide Ag2O, which melts at 300 C, but begins to decompose over the range 200 - 280.

Now, just speculating, but if AgNO3 decomposes to O2 (+ NO2), it may be thru 2 Ag2O --> 4 Ag + O2, that is, thru an intermediate Ag2O, which may be too difficult to isolate. However, this sort of reaction could be demonstrated by heating/melting AgNO3 with silver metal, and when the AgNO3 is all decomposed, the silver metal may have disintegrated, indicating a reaction where the metal was oxidized, then decomposed back to metal.

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  • $\begingroup$ Good point about dry chemistry ! $\endgroup$ – Poutnik Apr 3 '20 at 15:24

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