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Phosphate buffer is a common buffer in biological applications, it's especially popular for NMR. Magnesium ions are necessary for many biological systems to work, though the amount of magnesium ions you can add in a phosphate buffer is limited by the relatively low solubility of magnesium phosphate.

What I'm curious about is whether the phosphate buffer can affect the availability of magnesium ions even in concentrations where no visible precipitation occurs? Common phosphate buffer concentrations are around 10-100 mM, and magnesium ions are usually 5 mM or below. At those concentrations I see no visible precipitation, but is the magnesium actually fully available or is the phosphate sequestering it and thereby lowering the effective magnesium ion concentration in the solution?

Even just calculating the solubility of magnesium salts in phosphate buffer is not straightforward, as there are multiple species of magnesium phosphate and possibly even other salts like $\ce{MgKPO4}$.

Does phosphate buffer affect the availability of magnesium ions in solution compared to a non-interacting buffer? If that is so, is there a way to calculate the size of that effect?

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  • $\begingroup$ Could you provide a link to what phosphate buffer is and what it contains? $\endgroup$
    – F'x
    Commented Apr 29, 2012 at 18:23
  • $\begingroup$ @F'x Just potassium phosphate, not one of the fancier biological buffers. $\endgroup$
    – Mad Scientist
    Commented Apr 29, 2012 at 18:38
  • $\begingroup$ Technically, phosphate can and probably do form associates with magnesium ions. To find how strong, you should search databooks for relevant equilibrium constants, it is not such common info to keep it in mind if not needed. $\endgroup$
    – permeakra
    Commented Aug 25, 2012 at 10:21
  • $\begingroup$ I am using a buffer which has approximately 50 mM Phosphate and 7 mM Mg, There is no visible precipitate. But my experiments are Mg sensitive and at least 5 mM Mg is needed for proper response. My preliminary experiments are indicating that in the presence of phosphate free Mg is limited. $\endgroup$
    – user862
    Commented Nov 19, 2012 at 5:30

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If you're still interested in this, fluorescence may be a good method to test the phosphate dependence on magnesium activity. First record a fluorescence spectrum of a solution of magnesium and fluorescent dye, then add small volumes of concentrated phosphate and note any changes in the fluorescence. Of course, you would have to do a control experiment to see if the phosphate is directly perturbing the fluorescence. This write-up describes some magnesium-chelating dyes that you could use, though if you don't want to spend a lot of money on a dye, you may already have some chemicals in your lab that will work. Phenol looks like promising alternative as it fluoresces, is decently water soluble (83 g/L), and forms a complex with magnesium which breaks planarity, which would certainly change the fluorescence spectrum.

It would only take a day to get accurate quantitative results if you have access to a fluorometer.

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I don't have a source (at the moment) but it is accepted that high phosphate buffers do chelate/sequester Mg2+ ions. Pyrophosphate is particularly infamous for this. During RNA transcription experiments, the solution will get cloudy which is due exactly to this phenomena. It is also seen in cell-free protein synthesis which requires a steady concentration of Mg2+, that the increased concentration of the phosphate buffer requires an increased amount of Mg2+ to maintain the same effective free Mg2+ concentration [1].

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  • $\begingroup$ Do you have a source for plain phosphate? I have no pyrophosphate in the buffer. $\endgroup$
    – Mad Scientist
    Commented Apr 30, 2012 at 9:00
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If you deal with cell culture its irrelevant; because the magnesium channels located at the surface of the cell will be forcing magnesium ions to flow into the cell, and if the balance between free and chelated magnesium ions is disturbed in favor of chelate – it will produce free magnesium ions due to the chelate constant.

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  • $\begingroup$ I'm not dealing with any cells here, just simple, purified macromolecules $\endgroup$
    – Mad Scientist
    Commented Jul 20, 2012 at 19:14
  • $\begingroup$ So you should be more specific asking question – what do you mean by "availability of magnesium ions"? availability to what? if simple solution is involved you can calculate and compare the solubility ratio of your possible reaction $\endgroup$
    – java_xof
    Commented Jul 21, 2012 at 19:50

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