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I know how normal metal displacement reactions work. If one metal is more reactive than the other, they will displace.

For example, normally for this reaction: $$\ce{Fe_{(s)} + 2AgCl_{(aq)} -> Fe^{+2}_{(aq)} + 2Ag_{(s)} + 2Cl^{-}_{(aq)}}$$

I would say that because the iron is more reactive than silver, the iron pushes it's electrons onto the silver.

But in this case, can this reaction depicted actually occur? Because $\ce{AgCl}$ is an insoluble salt. This means that the $\ce{Ag+}$ ions are not floating around to touch the iron metal. This means the reaction cannot occur.

I think in reality since $\ce{AgCl}$ can dissolve a little bit due to equilibrium, the reaction would go ahead, but if we pretend equilibrium dissolution doesn't occur, can this reaction proceed?

I think it can occur if the iron is touching the AgCl directly (allowing for conduction). I also think it cannot proceed if the iron was not touching the AgCl. Is this correct? For the purposes of answering my question, please ignore the solubility of AgCl.

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    $\begingroup$ If one metal is more reactive than the other, they will displace. // While it is generally true, even less reactive metal ( slower reaction kinetics) can displace the more reactive metal from its salt, if it is preferred by the reaction thermodynamics by the sign of the Gibbs free reaction energy. // Insolubility is rather a vaque name convention for salts with very limited solubility. Reaction would be very slow, unless it is supported by dissolving of iron in acidic solution and by reduction with transient atomic hydrogen. $\endgroup$
    – Poutnik
    Mar 13, 2022 at 13:11
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    $\begingroup$ It doesn't matter if the reactants are salts or not: consider thermite, where aluminum displaces iron from its oxide. What would happen if AgCl, in contact with Fe, were warmed to 455 °C? $\endgroup$ Mar 13, 2022 at 21:37
  • $\begingroup$ States of aggregation should not be subscripted, it is not wrong, but the recommendations (Sec. 2.1.) are different. $\endgroup$ Mar 17, 2022 at 0:05

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In fact, at least on the basis of thermodynamics, a small amount of reaction between iron metal and "solid" silver chloride can be expected in water solvent.

This is because silver chloride is not completely insoluble in water. So if the thermodynamics is favorable, at least a small amount of reaction can occur.

First off, Ref. [1] quotes a standard reduction potential of -0.44 V for iron(II)/iron metal, versus +0.80 V for silver(I)/silver metal (all data and calculations are for 25°C and one atmosphere). Applying the Nernst Equation then gives the following equilibrium constant:

$\ce{2 Ag^+ + Fe <=> 2Ag + Fe^{2+}}, K = [\ce{Fe^{2+}}]/[\ce{Ag^+}]^2=8.5×10^{41}$

Taking silver chloride as having a solubility product of $1.77×10^{-10}$, we add in the dissolution reaction and render

$\ce{2 AgCl + Fe <=> 2Ag + Fe^{2+} + 2Cl^-}, K = [\ce{Fe^{2+}}][\ce{Cl^-}]^2=2.6×10^{22}$

This equilibrium condition is far beyond the range of concentrations/activities in a real aqueous ferrous chloride solution. So, what small amount of silver chloride dissolves into the water can undergo displacement favored by the large equilibrium constants above, producing ferrous chloride and silver metal, whereupon the removal of silver ion enables more silver chloride to dissolve. The reaction will eventually slow down due to accumulating chloride ions decreasing the solubility of the silver chloride, but the thermodynamics and mechanism are there for at least a little reaction to occur.

References

1. Veleva L., “Soils and Corrosion” (Chapter 32), in Corrosion Tests and Standards: Application and Interpretation, 2nd Edition, R. Baboian Ed., ISBN: 0-8031-2058-3, ASTM International, OH, pp.387-404, 2005.

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  • $\begingroup$ Thank you - but you haven't answered my question. I know that AgCl is not fully insoluable. Indeed I have referenced this in my original question, and how the slight dissolution would continually shift the equilibrium. My question was actually, what if we assume no Ag+ ions dissolved into solution? Would it still cause the deposition of silver onto the nail? $\endgroup$
    – John Hon
    Mar 19, 2022 at 11:50

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