Re: Magnesium precipitation

> From: Stephen.Pushak at saudan_HAC.COM
> Date: Fri, 29 Sep 95 1:16:23 PDT
> Subject: Re: Magnesium precipitation
> > Date: Wed, 27 Sep 1995 08:46:46 -0400 (EDT)
> > 	What reason do you have for saying that magnesium slowly precipitates
> > out of solution?  What does it come down as?  I can't see any reason to
> > chelate magnesium, are you sure you don't mean manganese (Mn)?
> I'm not certain about this. I recall reading someplace on a chelated
> supplement for gardens about chelated magnesium and about some magnesium
> compounds having a low solubility, but... I checked in my chemistry book
> and MgCO3 has solubility of 10**-15 (very low). I can't find any data
> on oxides of magnesium. It could be that what I recalled imperfectly was
> manganese? I'm no chemist so perhaps you can clear this up. BTW, that
> stinky gas that comes out of marshes; is that H2S? That stuff is very
> poisonous isn't it? Lethal in concentrations of ppm? One kind of swamp
> gas is methane, but I don't think that is the stinky stuff.

	The solubility of magnesium carbonate, according to the 
Handbook of Chemistry and Physics is  0.0106 g per 100mL.  This
would be a 0.00126 molar solution, implying a solubility product
of 1.59 x 10^-6.  That is, the product of the concentrations of 
magnesium and carbonate ions must be less than this.  Both concentrations
are expressed in gram-ions per litre (molar).  A gram-mole or gram-ion
of something is the molecular or ionic weight of it, expressed in grams.
e.g. 1 gram-mole of water weighs 18 grams (H20, H = 1 and 0 = 16).

	In aquarium water, the equilibrium

	HCO3-  <->  CO3--   +   H+

lies heavily to the left, because the equilibrium constant for the 
reaction is 4.84 x 10^-11, and [H+] is in the range 10^-5 to 10^-9
(pH 5 to 9).  Taking a specific example, at pH 7, the ratio of [HCO3-]
to [CO3--] is 2066.  Even at pH 9, it is 20.66.  The expression [X]
means the molar concentration of X, and the equilibrium constant for
the reaction above is defined as:

		-----------      (  =  4.84 x 10^-11)

	Equilibrium constant from Cotton and Wilkinson, Advanced
Inorganic Chemistry.

	Taking another specific example, if we have a carbonate hardness
of 6 degrees, there would be about 100 ppm clcium carbonate in the water.
However, we have just seen that there is very little carbonate in there,
it is almost all bicarbonate!  This solution is 10^-3 molar in calcium
(The molecular weight of CaCO3 is 100), and 2 x 10^-3 molar in bicarbonate.
The positive charges equal the negative ones, and all the other charged
species (H+, OH- and CO3--) are in considerably lower concentrations in
the pH range of interest to us.
	In fact, in this case at pH 7, [CO3--] is about 10^-6 (molar), so
the required value of [Mg++] for precipitation of the carbonate would be
1.59 x 10^-6 divided by 10^-6, or 1.59 molar.  This is a huge concentration,
over 38 grams per litre of magnesium (38000 ppm).  At pH 9 you could still
get 380 ppm before MgCO3 could come down.

	I hope that what I have written here is clear.  I'm in the middle
of writing up an explanation of water hardness, pH and CO2 concentrations
which I shall put either here or in the .aquaria newsgroups.  The point is
that if you want to make sense of chemical equilibria, you do have to
understand some chemistry and watch your units when looking at concentrations.

	The "stinky gas" would have some H2S in it, and the H2S
would be responsible for most of the stink.  H2S is lethal to humans if the
concentration is much above about 20 ppm (volume concentration in air), though
the time taken to kill is very dependent on the concentration, and it is
very dangerous at least in part because one's nose stops smelling it when
the concentration is high.

Paul Sears        Ottawa, Canada.