If one were to dissolve magnesium citrate (about $100~\mathrm{mg}$) in water (about $70~\mathrm{ml}$) and let the solution sit for 48 hours, would the solution be exactly the same as one dissolved and consumed right away?

How does one explain the difference in taste from a 48-hour solution and a 5-minute solution?


The only reactions that I would expect to occur to any significant extent with pure water are acid-base reactions between the citrate and water, where water, being amphiprotic, could act as either an acid or a base:

$$\ce{C6H6O7^2- + H2O <=> C6H7O7- + OH-}$$ $$\ce{C6H6O7^2- + H2O <=> C6H5O7^3- + H3O+}$$

Citrate, being the conjugate base of a weak acid, is itself a weak base (according to the citric acid Wikipedia page, the $\mathrm pK_\mathrm a$ for the dissociation of the second proton is 4.75, and for the third between 5.41 and 6.40). Citrate is tripotic, and in the magnesium salt citrate possesses one undissociated acidic proton, which makes it an amphiprotic molecule as well. Given that the $\mathrm pK_\mathrm a$ of water is approximately 15.7, and the $\mathrm pK_\mathrm a$ of hydronium ($\ce{H3O+}$) is −1.74, both reactions above favor the reactants over the products by many orders of magnitude, though the second reaction is favored above the first (i.e., the dibasic magnesium citrate salt should still be very slightly acidic, overall).

Acid-base reactions typically reach equilibrium quite quickly, so there should be very little change in pH and no pH-attributable differences in taste over time. If any change does occur, I would expect it to be the result of a change in concentration as the water gradually evaporates. Another possibility is that, if the water is impure (which tap water certainly is), additional side-reactions might occur. What those are would depend on the exact composition of the water as well as environmental conditions. It's possible for the magnesium to form a precipitate, for example, in combination with other anions dissolved in the water (unlikely given the quantity in your example, but possible). If the magnesium citrate is part of a commercial product containing other ingredients, then it may well be that other agents are responsible for any effects you observe.


The difference in taste is due to a difference in molarity of the magnesium citrate.

Its solubility in water is 20 g/100ml (http://en.wikipedia.org/wiki/Magnesium_citrate)

Over a 48-hour period a significant amount of water will evaporate from the container meaning the concentration will increase. Based on the numbers provided it is unlikely that the concentration will get the the point of precipitation; the solution will visually be the same.

Evaporation is dependent on surface area so the larger the surface area of exposed water the more profound the difference in taste will be.

  • $\begingroup$ Would keeping the container closed (to avoid evaporation) still have an effect on taste? $\endgroup$ – mythealias Jul 24 '13 at 13:30
  • $\begingroup$ Following the logic presented in my answer yes. Preventing evaporation would prevent a change in the molarity of the solution. It there is some type of chemical decomposition of an actual molecule covering the solution would have no effect. $\endgroup$ – anglinb Jul 28 '13 at 18:18

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