Sciencemadness Discussion Board

MgSO4 + citric acid + ? -> Food grade magnesium citrate

CrimpJiggler - 30-3-2014 at 05:07

MgSO4 and citric acid are used in large quantities (i.e. MgSO4 as a bath salt) so they're pretty cheap, but supplements containing magnesium citrate or oxide are too expensive for me to use daily, since magnesium isn't the only supplement I take. Mg(OH)2 would be the ideal starting material for making the citrate salt, but milk of magnesia is too expensive. Citrate can chelate the Mg2+ ions, and although it can't wrap around it like longer chained chelators could (i.e. the way pentetic acid can completely encage ions metal ions), it can form complexes (3:2) complexes like this:

which are much less water soluble (and should be easy to crash out of an aqueous solution).

Two things I'm not sure about though are what solvent would be best, and what to do about the increase in pH that'll occur when citric acid loses its protons. The latter is the more important question cuz if the protons are gonna be pairing up with sulphate counterions, then I'd end up with a concentrated H2SO4 which will start causing side reactions. I'm thinking NaHCO3 or Na2CO3, adding it in small amounts whenever the pH drops below 6. Then lets say my solvent is glycerol, sodium sulphate will crash right out, but with all that potential for hydrogen bonding, I reckon the citrate complex will remain dissolved. And if you were to make the solution sufficiently concentrated, then there might not be any need to isolate the citrate complex, you could just drink a small volume of the solution.

blogfast25 - 30-3-2014 at 05:21

It's unlikely that you can prepare magnesium citrate this way.

You will need to convert the magnesium sulphate to magnesium carbonate or hydroxide, filter it off and wash it, then react it with a solution of citric acid. I seem to recall there's a thread on this already: try using the search facility.

There are other things that are plain wrong with your reasoning but I'll leave it to others to elaborate.

[Edited on 30-3-2014 by blogfast25]

CrimpJiggler - 30-3-2014 at 10:32

I'll see if I can spot the flaws in my reasoning you're referring to. What I said about pentetic acid applies to huge paramagnetic f-block metals like gadolinium, magnesium would probably be way too small to form one of these caged complexes:
that complex is used as an MRI contrasting agent for brain imaging, its lipophilic enough to readily cross the blood brain barrier.

Other than that I can't spot anything. As for going through MgCO3 or Mg(OH)2, I don't see how that would be any easier. Well, the hydroxide is pretty insoluble in water so I suppose I could make that easily enough with KOH or the likes. I haven't a clue how I would make the carbonate but I don't see how these metathesis reactions are any more feasible that what I proposed.

BTW I saw that thread. I assumed it irrelevant since the OP is using MgO as a starting material. In that case, the Mg would be tightly bound to the oxygen atom, I bet it has really bad water solubility. With MgSO4 on the other hand, the Mg2+ ions dissociate with ease.

We'll find out if magnesium citrate can be made like this soon enough, just need to get some epsom salts, have everything else.

[Edited on 30-3-2014 by CrimpJiggler]

blogfast25 - 30-3-2014 at 11:58

For one, what you call a 'complex' is really just a salt. According to Wiki, it has a solubility of about 20 g / 100 ml water. When you mix magnesium sulphate solutions and citric acid solutions you have just that: Mg cations, sulphate anions and weakly dissociated citric acid in water.

"[...] and what to do about the increase in pH that'll occur when citric acid loses its protons"

When an acid deprotonates in water, H<sub>3</sub>O<sup>+</sup> ions are formed and the pH goes DOWN, not up.

Glycerol... do you realise just how viscous that is?

No. Take the required amount of Epsom salt, dissolve it in water and add the required amount of NH3 solution, acc.:

Mg2+ + 2 NH3 + 2 H2O === > Mg(OH)2 + (NH4)2SO4. Filter off the Mg(OH)2 and wash out the ammonium sulphate. Then dissolve the Mg(OH)2 in the required amount of citric acid solution. Now you have at least a sulphate-free magnesium citrate solution.

[Edited on 30-3-2014 by blogfast25]

AJKOER - 31-3-2014 at 18:06

Since you are interested in food grade products, I would recommend the following path to MgCO3. Add an aqueous solution of MgSO4 to aqueous NaHCO3. Heat to decompose the Mg(HCO3)2 to insoluble MgCO3 and CO2. Then, add the washed MgCO3 (which is also used to treat acid indigestion) to citric acid.

I would not add citric acid slowly to Mg(OH)2 as this could form some insoluble basic magnesium citrate.

[Edited on 1-4-2014 by AJKOER]

CrimpJiggler - 2-4-2014 at 12:05

AJKOER: Thanks. Sounds like a brilliant route.

blogfast25: Cheers for elaborating. Well, everything besides the pH part, that was a typo, I know how the pH scale works. I suppose the term complex is reserved for compounds where a metal atom is binded to one or more chelating agents via dative bonds. Thats a coordination complex at least, I don't know if theres a generic term "complex" in chemistry or what.

I'm familiar with glycerols viscosity but I didn't consider how difficult it would be to run it through a filter in order to collect the precipitate. I bet you could thin it out with ethanol or IPA though without altering its solvent properties too much. I'll test this out and in a few minutes. I've experimented with adding menthol to e-cigarette liquid, its pretty insoluble in propylene glycol so first I dissolve it in acetone or IPA then mix with the PG, shake, and allow the mixture to homogenize. It thins out the propylene glycol significantly, without causing anything to crash out, so I think the same would work for glycerol.

I didn't know that about NH3, thanks.

Both pretty cool routes, thanks a lot!

[Edited on 2-4-2014 by CrimpJiggler]

blogfast25 - 4-4-2014 at 04:38

Good luck.

CrimpJiggler - 5-4-2014 at 08:44

As a quick test, I mixed some warm KOH solution, to some warm MgSO4 solution and observed the liquid turn cloudy white as the Mg(OH)2 precipitated out. It formed a suspension, I doubt I could filter that out of the liquid very easily. I then tested out Maybe with very fine filter paper and a decent vacuum filtration setup, but I don't have either of those. I'm gonna see how MgCO3 interacts with water next. I'll try out the NH3 route when I have more free time. That should work for Ca(OH)2 too, although food grade Ca(OH)2 is probably cheaper than epsom salts themselves so I'll only be doing that for the sake of knowledge.

CrimpJiggler - 5-4-2014 at 09:20

MgCO3 forms a similar milky suspension, IME Na2SO4 crystallizes out so I don't think Na2SO4 has much to do with this suspension. Maybe it won't be so hard to filter out. Would it be best to use a saturated MgSO4 solution? Also, what advantages does using NaHCO3 have over using Na2CO3?

blogfast25 - 6-4-2014 at 03:58

Quote: Originally posted by CrimpJiggler  
MgCO3 forms a similar milky suspension, IME Na2SO4 crystallizes out so I don't think Na2SO4 has much to do with this suspension. Maybe it won't be so hard to filter out. Would it be best to use a saturated MgSO4 solution? Also, what advantages does using NaHCO3 have over using Na2CO3?

Even if you chill the solution to 0 C Na2SO4 still has a residual solubility of about 4.9 g/100 g water (Wiki solubility table), so you won't get rid of all of it.

Re. Mg(OH)2/ 'carbonate' suspensions, try just standing them for a few days: the grains tend to grow, thus making them easier to filter.

In NaHCO3 (which is only sparingly soluble) the actual carbonate (CO<sub>3</sub><sup>2-</sup> species) is fairly low compared to the OH<sub>-</sub> concentration. As a results in several cases bicarbonate has a tendency to precipitate either a hydroxide or a basic carbonate. Magnesium bicarbonate is fairly soluble. AJ is right though that boiling a solution of Mg(HCO3)2 should drive off the CO2 and precipitate the carbonate:

Mg(HCO3)2(aq) === > MgCO3(s) + CO2(g) + H2O(l)

You can remove the last bit of CO2 by shaking the cooled suspension (as you would with a pop bottle to end the fizz).

By contrast a (say) sodium carbonate solution contains more CO<sub>3</sub><sup>2-</sup> ions than OH<sup>-</sup>, so it has more of a chance to precipitate a pure carbonate.

With precipitated carbonates, making sure you have a well defined structure (like MCO<sub>3</sub>;) rather than an ill-defined 'basic carbonate' is harder to achieve. And with some cations (like aluminium) it doesn't even matter whether you use sodium carbonate or sodium bicarbonate: what precipitates is Al(OH)3!

Precipitating compounds from saturated solutions does have one disadvantage: as a general rule more occlusion will occur. This means that your precipitate will contain more soluble salts, locked into the solid (and thus hard to wash out). Purer precipitates are obtained at lower concentrations, as a broad rule.

[Edited on 6-4-2014 by blogfast25]