Magnesium is one of the most abundant elements in the Universe and on Earth....

So why is it not used in the forms of Magnesium Carbonate and Magnesium Phosphates (Hydroxyapatites, etc.) to make shells and bones?

Is it not as widely available in the oceans and/or crust?

Is it not quite as 'versatile', chemically speaking, as calcium?

Is it not soluble enough, or too soluble?

Or are Magnesium compounds just too brittle or weak?

I've always wondered about this.....

  • 1
    $\begingroup$ This question might be better suited to SE Chemistry, using the biochemistry tag. $\endgroup$
    – Fred
    Jan 11 at 5:58
  • $\begingroup$ Also, even though magnesium may be more abundant than calcium doesn't mean it is as readily available as calcium is, particularly for biological purposes. The way that it is incorporated into igneous minerals may mean it may be more difficult for biological processes to access. $\endgroup$
    – Fred
    Jan 11 at 18:28
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    $\begingroup$ Calcium is much more common in the Earth's crust than is magnesium, almost by a factor of two. $\endgroup$ Jan 13 at 12:03
  • $\begingroup$ This is complicated by the fact that you need dissolved magnesium in order to form aragonite, which is currently the most common calcareous mineral in shells and bones. $\endgroup$
    – Spencer
    Jan 15 at 0:39

1 Answer 1


First off, magnesium is not more common than calcium where it counts for Earth organisms, namely the crust. By mass calcium is more common in the crust, by atoms they are about equal (the enrichment of calcium in Earth's crust is described here). Plus, at least some organisms might prefer to save their magnesium for chlorophyll, which is rather important for biology in Earth in its own right.

Calcium and magnesium are both alkaline earth metals, but their compounds have subtly different properties. Solubility, which us a major consideration for forming stable and robust skeletal structures, is among them. Among carbonates, calcium carbonate has a solubility of $0.013\text{ g/L}$ at $25°\text{C}$; the corresponding figure for magnesium carbonate is $0.139\text{ g/L}$. When magnesium carbonate is dissolved in water, much of it is in the form of ion pairs which favors greater solubility and also makes that solubility greater than that predicted from the usual solubility product. Calcium ions, being bulkier, are less effective at forming ion pairs with multiple charged anions. Clearly calcium carbonate is more likely to precipitate from ocean water to form shells, pearls, etc. Even where there is magnesium carbonate, it is often in the form of a double salt with the less soluble calcium carbonate component (dolomite).

Phosphates tell a similar story. Wikipedia does not give quantitative data for the solubility if $\ce{Mg3(PO4)2}$ or $\ce{Ca3(PO4)2}$, but the latter has a lower solubility product ($2×10^{-29}\text{(Ca)}<1×10^{-25}\text{(Mg)}$) and again, we would expect ion pairing with a multiple charged anion to favor magnesium phosphate dissolution even more. So as with carbonates, calcium phosphate is easier to precipitate and more resistant to (re)dissolution than the magnesium counterpart.


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