is there iron rich red clay - The Planted Tank Forum
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post #1 of 7 (permalink) Old 11-19-2015, 10:19 PM Thread Starter
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is there iron rich red clay

iron rich red clay, can it be used as iron tabs?
I see some people using Amaco Mexican Pottery clay

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post #2 of 7 (permalink) Old 11-20-2015, 12:30 AM
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yeah, most likely. You can add other micros AND macros into the clay too.

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post #3 of 7 (permalink) Old 11-20-2015, 03:03 AM
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I used API laterite as a mix in with my MTS. It's kind of a mess now on my black sand, but it's kept my AR mini a beautiful red. Some red clays have aluminum mixed in, which can be dangerous in an aquarium.
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post #4 of 7 (permalink) Old 11-20-2015, 03:40 AM
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Plants take up nutrients as ions. For example KNO3 becomes K+ and NO3- when dissolved in water, and the plants take up both ions for food. Rocks and clay tend to be minerals which are not water soluble, so their constituents are not in ion form for the plants to be able to use them. Bacteria can slowly convert those minerals to ion forms to make the components available, but that is not a quick process.

Yes, some clays do contain iron oxides, but the iron doesn't become available as ions when used in a substrate: iron oxides don't dissolve easily in water. The iron we dose as a plant food is in forms that do dissolve in water and can be available to the plants, especially if combined with a chelator which protects the iron ions from quickly oxidizing into unavailable molecules.

You can tell that I am not a chemist, but I believe I am right on this even if the details are a bit fuzzy.

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post #5 of 7 (permalink) Old 11-20-2015, 08:18 AM
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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.
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post #6 of 7 (permalink) Old 11-20-2015, 02:47 PM
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I once expierimented with the red pottery clay and made some little plug's filled with Osmocote and then baked them on cookie sheet in oven and after they were good and dry,I inserted them into the substrate convinced I had done a good deed for my plant's.
Was not until six month's had passed and I tore down the tank thinking my soil had give all it could with respect to nutrient's,that I discovered the little plug's I made were still intact and prolly provided very little in the way of nutrient's.
They were just wet,hard,dirty.
I now just freeze some osmocote pellet's in ice cube tray,and use these for root tab's.
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post #7 of 7 (permalink) Old 11-20-2015, 06:07 PM Thread Starter
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Thank you for that lesson, great explanation!
Will go with the recommended route
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