Re: Iron ions
To: Aquatic-Plants at actwin_com
Subject: Re: Iron ions
From: psears at NRCan_gc.ca (Paul Sears)
Date: Wed, 30 Oct 1996 09:01:58 -0500 (EST)
In-Reply-To: <199610300839.DAA15728 at looney_actwin.com> from "Aquatic-Plants-Owner at ActWin_com" at Oct 30, 96 03:39:01 am
> From: Marque Crozman <marquec at gastech_com.au>
> Subject: Fe2+ and Fe3+ ions
> > Can aquatic plants really not cope with FeIII in the substrate?
> >The chelated iron we add is a _very_ tightly bound FeIII complex with
> >EDTA, and they use that, it seems. Why can they not use FeIII in solid
> >material in the substrate, provided the roots can reach it? _Must_ the
> >substrate go anoxic?
> According to Gesting (Berti Gesting-"Nature and Aquarium-Advanced Aquarium
> management-aquarium plants")
I'm afraid I have given up trusting aquarium books on anything
to do with chemistry. So much of what is written in them is confused,
half-right, misleading or flatly wrong.
> "Some ions can, of course, be present in a different form, eg Iron as Fe3+
> of Fe2+ and manganese as Mn4+ instead of Mn2+, but it means that they are
> longer plant accessible.
Fe3+ ions do precipitate out of anything other than very acidic
solution, and the resulting solids will clearly be useless to plants if
they are in the filter. My question concerns FeIII compounds (solid ones)
_in_ the substrate. Is the iron usable by the plants if the roots can reach
the solid? Plants manage very impressive oxidation state changes in other
cases, e.g., nitrate -> nitrogen in reduced states, e.g., amino acids.
In the case of manganese, if you put the Mn in as a Mn2+ salt, its
oxidation to Mn(IV) is highly unlikely - there are very few stable Mn(IV)
compounds. In the case of MnO2, it is readily attacked by reducing agents
in acidic solution(*), so I would expect plants to be able to deal with that
if it were in the substrate.
* Cotton and Wilkinson, Advanced Inorganic Chemistry.
> On another rusty topic:
> Something that I cant get out of my mind was a talk given, where the
> units of ppm and mg/l were tauted as being approximately the same.
They are. We are talking about ppm _by weight_.
A litre of water weighs very close to 1 kg, so if you dissolve
1 mg of something in 1 L of water, you have 1 mg in 1 kg, which
is 1 part per million.
> Avogadro's Number, 6.022E23 using scientific notation, is the number of
> molecules that would be present in 1 mole of any substance. As Fe ions are
> present as single atoms in solution, 1 mole of Fe ions would weigh 55.8g
> there would be 6.022E23 atoms.
> H2O then has an atomic weight of 17.9994 or about 18
> So one mole of pure H20 has a weight of 18grams. Therefore one litre of
> (which weighs 1kg) contains 55.56 moles or (55.56 * 6.022E23) or 3.35E25
> molecules of H2O.
> To create a solution then, that contained 1mg/l of Fe ions, we would have
> to add
> (6.022E23 / 55.8) or 1.08E22 ions of Fe to our 1 litre of water.
No. This is the number of ions in 1 g of iron, not 1 mg. The
number you want is one thousandth of the one you got: 1.08E19
> We would than have 1.08E22 ions of Fe in 3.35E25 molecules of H20. Or to
> it another way,
> 322.4 ions of Fe for every 1millon parts of Fe ions and water molecules.
> ie: 1mg/l = 322.4 ppm.
I make it 1 mg/L is 0.3224E-6 iron ions per water molecule.
Incidentally, Avogadro's number goes out of the calculation, all you
needed was the weight ratio (10E-6) and the ratio of atomic/molecular
weights (18/55.8 = 0.3226).
This result (0.3226E-6) is the _number_ ratio. We actually _use_
the weight ratio.
Paul Sears Ottawa, Canada
Finger ap626 at freenet_carleton.ca for PGP public key.