Why is AgCl most soluble in water?

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I heard that:

Silver chloride is more soluble in very concentrated NaCl solution (brine) than in pure water.

Is it true? I mean, won't the common ion effect operate? Or is this maybe due to some kind of complex formation as the NaCl solution is highly concentrated?

asked Oct 7, 2018 at 4:46

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I mean won't the common ion effect operate?

The common ion effect always applies for pure substances. With the added sodium chloride, you have changed the effective solvent and changed the ions.

Or may be this is due to some kind of complex formation as NaCl solution is highly concentrated?

Exactly, before your ionic equation was: $$\ce{Ag+ (aq) + Cl- (aq) <=> AgCl (s)}$$ but with brine, the ionic equation is: $$\ce{AgCl (s) + Cl- (aq) <=> [AgCl2]- (aq)}$$ Which has a higher solubility constant, and where as before more chloride made the $\ce{AgCl}$ less soluble, if you note that you have switched regimes and $\ce{Cl-}$ has switched from precipitating a silver ion to dissolving a silver chloride molecule, which means now with increased chloride concentration, silver chloride becomes more soluble.

answered Oct 7, 2018 at 22:22

A.K.A.K.

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The answer by A.K. is eminently reasonable, but I wondered if stability constants were available. I quote from the summary of Ref. 1 :

The solubility of silver chloride in concentrated solutions of various chlorides was determined mainly at 25°. The solubility is nearly doubled in going from o° to 25°, the rate of increase above and below 25° being nearly logarithmic. The solubility is sharply proportional to integral powers of the chloride concentration throughout considerable ranges, a fact explained by assuming the existence of the complex anions AgCl3$^{^-2}$, AgCl4$^{-3}$ and possibly AgCl5$^{-4}$. No evidence of the ion AgCl2$^{^-}$is found.

It seems that the simplest explanation is only the beginning of understanding the solubility increase.

References

1. The Solubility of silver chloride in chloride solutions and the existence of complex argentichloride ions, J. Am. Chem. Soc. 1911, 33, 12, 1937– 1946 Publ. Date: December 1, 1911 https://doi.org/10.1021/ja02225a008

Buck Thorn♦

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answered Jul 7, 2021 at 17:17

James GaidisJames Gaidis

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REFERENCES

1. Wagman, D. D.; Evans, W. H.; Parker, V. B.; Schumm, R. H.; Halow, I.; Bailey, S. M.; Churney, K. L.; Nuttall, R. L. The NBS Tables of Chemical Thermodynamic Properties: Selected Values for Inorganic and C1 and C2 Organic Substances in SI Units, J. Phys. Chem. Ref. Data 1982, 9, Suppliment 2.

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2. Lias, S. G.; Bartmess, J. E.; Liebman, J. F.; Holmes, J. L.; Levin, R. D.; Mallard, W. G. J. Phys. Chem. Ref. Data, 1988, 9, Suppliment 1.

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3. While there are more contemporaneous techniques to deriving enthalpies of formation of solid salts and lattice energies, for the current study it suffices to consider the careful and thorough review by Hisham, M. W. M., Benson, S. W. In From Atoms to Polymers: Isoelectronic Reasoning; (Liebman, J. F.; Greenberg, A. eds.) VCH: New York, 1989.

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4. Chase, Jr., M. W. NIST-JANAF Thermochemical Tables 4th ed.; J. Phys. Chem. Ref. Data Monogr. 9, 1998.

5. As may be deduced from ionic radius ratios or directly found from the textbook literature, one can not even “blame” the crystal lattice type, e.g. see Huheey, J. E.; Keiter, E. A.; Keiter, R. L., Inorganic Chemistry: Principles of Structure and Reactivity, 4th ed.; HarperCollins, New York, 1993. One finds that NaCl, KCl, and AgCl all have the same 6-coordination for the metals and chlorine alike.

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6. Fox, B. S.; Beyer, M. K.; Bondebey, V. E. J. Amer. Chem. Soc. 2002, 124, 13613. Unfortunately, there is seemingly no complementary combined experimental/theoretical study involving the corresponding Na+-or K+-containing complexes known to the current author.

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7. Something is wrong here. From textbook knowledge of solubility behavior, we know that NH3 solubilizes AgCl, although the exothermicity and exoergonicity of reaction (9b) suggests the opposite. We can only conclude that the key silver–ammonia ion in solution is not [Ag(NH3)2]]+ but rather some more highly ammoniated species, such as the hydrogen bonded complex, [Ag(NH3)2]+ • NH3, a conclusion consistent with the calculational results of Fox, Beyer, and Bondebey, Ref. [6].

8. Solutions of sodium in liquid ammonia just narrowly escape seeming irrelevance as well. Stated differently, from Ref. [1], we find that the enthalpy of formation of such solutions range from −6 kJ-mol−1 for 1:40 solutions to −2 kJ-mol−1 for 1:10000 solution. Such negative values attest to considerable stabilization—after all, one must overcome the 107 kJ-mol−1 needed to atomize, i.e., sublime, the solid metal. However, it is scarcely enough to allow for the favorable exothermicity (and presumably exoergicity) needed to form a solution of dissolved metal. We find, admittedly unfortunately, that there is seemingly no complementary data involving related solutions containing potassium or silver.

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