That's right Hoppy.
Simplified Pourbaix diagram for 1 M iron solutions.
This shows what state iron is in dependent on pH and ORP.
The solid lines represent where the chemical species are at equilibrium. Line a
shows the pH at which half of the 1 M iron is Fe+++ and half is precipitated as Fe(OH)2.
At pH 7.0 and ORP +400mv, any free iron in the water will want to be Iron(III) oxide-hydroxide
( Fe(OH)3 ), which is a solid. As long as that water stays around that pH and that ORP, it will remain as a solid and unavailable to plants.
At pH 2.0 and ORP +200mv, any free iron in the water will want to be free Fe++ cations, completely dissolved in the water for direct use by plants, either in life or in death.
In a substrate, generally two things happen, pH drops and oxygen level (ORP) also drops, which will mean that any precipitated (solid) iron will begin to dissolve into free Fe++ cations. As water flow is conductive to erosion, as is water flow to dissolution rate. So while a substrate generally has pH and ORP in its favor, it generally does not have water flow in it's favor. So while precipitated iron will dissolve in the substrate, it will be at a slow dissolution rate, however, the store of free Fe++ will gradually increase with time
. Some of it will stay in the substrate as free Fe++ cations, some of it will be uptaken by plant roots, some of it will escape to the water column. But some of the free Fe++ cations that escape to the water column will again precipitate in the favorable pH and ORP levels of the water column, to form a solid, sink to the substrate and continue the cycle. If the dissolution rate of the solid form of iron is greater then plant uptake via roots, it's concentration will increase.
While free Fe++ is great for it's ease of use (from a plants perspective), it's also the most toxic form of the element. In the substrate, only some percentage will be attached to plant roots, and some percentage of that will be uptaken by the plant. Since flow is non-existent, most of it will simply be making it's way slowly through the substrate.
So it's probably one of the most safest methods to dose an aquarium with iron, since only a small percentage will be in the water column, or attached to plant roots, provided
, excess precipitated iron that's sitting on the substrate, or plant leaves, or otherwise in direct contact with the water column is regularly removed. If the water column parameters (pH and ORP) change to favorable free Fe++ conditions, then the water column has all three conditions suitable for the transformation of precipitated iron to free Fe++ cations. Ph, ORP and flow
. Imagine spilling an entire 500ml solution bottle of iron into your aquarium, only without the chelated protection.
Chelation is a complex bond of other atoms with iron to form the protection. Iron sulfate (FeSO4) is a simple Fe (cation) and SO4 (anion) bond and easily becomes precipitated iron ( Fe(OH)2 ), a Fe (cation) and two OH (anion) bond. EDTA Fe
is a bond of ten Carbon atoms, thirteen Hydrogen atoms, the Fe atom, two Nitrogen atoms and eight oxygen atoms, with DTPA Fe being even more complex. The protection of the chelation provides a slow release of free Fe+++, meaning that some will be uptaken by plants, with the remaining being precipitated to Fe(OH)2 over time. However, the protection provided by chelation doesn't protect against user error. Add chelated iron at a faster rate then plant uptake and the concentration of precipitated iron will increase.