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Iron: The most problematic nutrient

Its behavior in slightly acidic, low-buffered, low-organic-matter, high-oxygen freshwaters (i.e. in planted aquariums)

System parameters

  • Slightly acidic water → pH ≈ 5
  • High oxygen levels → redox potential (Eh) around +300 to +400 mV
  • Low bicarbonate levels → poorly buffered
  • Low organic matter levels → low dissolved organic carbon (DOC)
  • Medium temperature → ~25°C (77°F)
  • No substrate → for better control
  • Moderate [water] circulation → for better nutrient distribution
  • Slightly increased CO2 → 5-10 mg/ℓ

Outline

This is the result of a detailed research task to provide a comprehensive breakdown of iron behavior in slightly acidic, low-bicarbonate, low-organic-matter, high-oxygen freshwaters, focusing on submerged aquatic plants.

Aspects:

  • iron chemistry
  • speciation
  • reactions (and their speed → kinetics)
  • plant uptake mechanisms
  • toxicity risks
  • unique behavior of oligotrophic (soft-water) plant species

Key aspects of iron chemistry in such environments

First, iron speciation is crucial. At pH ~5, Fe3+ is more soluble than at higher pH, but it still hydrolyzes. Redox potential around +300 to +400 mV is oxidizing. In such conditions, Fe3+ dominates over Fe2+. But even then, Fe3+ forms insoluble hydroxides. So, dissolved iron would mainly be as Fe3+-organic complexes (Fe3+-OM).

Forms available to plants: probably dissolved Fe3+ complexes and possibly colloidal forms, but plants might reduce Fe3+ at the root surface via reductases.

What if we add 0.5 mg/ℓ of Fe as FeCl3 or FeSO4 to this system? Adding FeCl3 would lead to rapid hydrolysis and precipitation as ferrihydrite. The kinetics would be fast, seconds to minutes. FeSO4 would oxidize to Fe3+, but in oxygen-rich water, oxidation is quick, especially at higher pH. However, at pH 5, oxidation is slower. The presence of organic matter could slow precipitation by forming complexes. So, the added iron might end up as particles or adsorbed to organics.

Toxicity is another concern. Free Fe2+ can cause Fenton reactions, producing ROS. But in high-oxygen, low-iron water, adding iron might increase ROS risk. Plants have antioxidants, but sudden exposure could stress them. Also, precipitation might reduce bioavailability, leading to deficiency if iron is sequestered.

We should strive for a balance between iron's necessity and toxicity. In oligotrophic systems, even small additions can shift dynamics. Since iron is a trace element, the line between deficiency, sufficiency, and toxicity can be very thin.

Comprehensive breakdown of iron behavior

1. Iron Uptake Partitioning Estimates

Most aquatic plants absorb iron through their roots from sediment. Only plants with minimal root systems (typically stem plants, for example) absorb most of their iron through their leaves from the water column. Plants absorb most nutrients (including iron) through so-called transporters.

What Is a “Transporter”?

Imagine each plant cell as a tiny, water-tight house. The cell wall is like a picket fence and the cell membrane is the front door. Outside the house is the “yard” of water or soil that contains food—‐minerals such as iron, nitrogen, or potassium—that the plant must bring inside to stay alive.

The trouble is that these minerals can’t simply drift through the closed door. This is where a transporter comes in.

A Simple Analogy

Think of a transporter as a specialized doorman or gate built into the front door:

  1. Recognition key
    • The doorman is trained to recognize one-particular nutrient—for example, iron.
  2. Active carry-in
    • When that nutrient shows up at the doorstep, the doorman grabs it, pulls it through the doorway, and releases it inside.
  3. Selectivity
    • Each doorman works only for its assigned cargo; the iron doorman ignores potassium, and vice-versa.
  4. Energy use
    • Sometimes the doorman even uses the plant’s energy (like ATP or a built-in battery of charged ions) to haul nutrients in against the flow.

Why Transporters Matter

  • Survival in poor conditions – In nutrient-poor lakes or soils, transporters with “extra-sticky hands” (high affinity) let plants snatch scarce minerals before they wash away.
  • Balanced diet – Different transporters handle dozens of nutrients, preventing shortages or toxic overloads.
  • Adaptation – Aquatic plants that live floating in water shift many doormen to their leaves, while rooted plants station most at their roots.
    Where in the plant (roots vs leaves) are most of the iron transporters located:
    Plant TypeConditionRootsLeavesNotes
    Plants with
    Extensive roots
    • Vallisneria
    • Echinodorus
    • Potamogeton
    • Sagittaria
    • Eleocharis
    		Sediment-anchored95-100%0-5%
    • Roots access sediment Fe2+/Fe3+-OM
    • Rhizosphere acidification/reduction
    • Leaves acquire colloidal Fe
    Floating roots
    (no sediment)    		↓↑
    • Roots absorb dissolved Fe2+/chelates
    • Exudate dilution reduces root efficiency
    • Leaves capture Fe(OH)3 colloids
    Plants with
    Minimal roots
    (Stem plants)
    • Rotala
    • Cabomba
    • Egeria
    • Myriophyllum
    		Sediment-anchored0-10%90-100%
    • Roots less developed; rely on sediment Fe2+
    • High leaf:root ratio favors foliar uptake
    • Leaves reduce Fe3+-colloids
    Floating roots
    (no sediment)    		↓↑
    • Roots inefficient without anchorage
    • Leaves critical for Fe(OH)3 colloid capture
    • Species like Cabomba absorb >50% via leaves

Summary

  • Transporters are protein gates in cell membranes.
  • Each one is highly specific for a certain nutrient.
  • They actively move nutrients from the outside world into the cell, keeping the plant healthy.

Table 1: Dominant High-Affinity Transporters in Rhizomatous/Rosette Aquatic Plants with Extensive Roots

  • e.g., Vallisneria, Echinodorus, Potamogeton, Sagittaria, Eleocharis
Nutrient Primary
Transporter		 Root % Leaf % Notes Km
(μM)		 Km
(mg/ℓ)
NO₃⁻ NRT1 90-95% 5-10% Leaf uptake in N-rich waters 10-30 0.6-1.9
NH₄⁺ AMT1 95-98% 2-5% Trace in leaves 5-20 0.09-0.4
PO₄³⁻ PHT1 95-100% 0-5% Leaves lack high-affinity P transporters 1-10 0.1-1.0
K⁺ (high) HAK5 95-98% 2-5% Root-specific high-affinity uptake; essential in low-K waters 2-20 0.08-0.8
K⁺ (low) AKT1 30% 70% Low-affinity uptake; dominant in leaves when [K⁺] > 0.5 mM 500-1000 20-40
Fe²⁺ IRT1 95-99% 1-5% Root-exclusive in rooted species 0.1-0.5 0.006-0.03
Fe³⁺ FRO2 → IRT1 100% 0% Reduction occurs only at roots n/a n/a
Zn²⁺ ZIP3 80-90% 10-20% ↑ in leaves of rootless species 0.3-2.0 0.02-0.13
Cu²⁺ COPT1 85-90% 10-15% Leaf role in photosynthesis 0.01-0.1 0.001-0.006

Table 2: Dominant High-Affinity Transporters in Stem Plants with Minimal Roots

  • e.g., Cabomba, Egeria, Myriophyllum, Rotala
Nutrient Primary
Transporter		 Root % Leaf % Key Adaptations Km
(μM)		 Km
(mg/ℓ)
NO₃⁻ NRT1.1 (Leaf) 10-20% 80-90% NRT1.1/NPF6.3 upregulated in leaf epidermis/veins 10-20 0.6-1.2
NH₄⁺ AMT1;3 (Leaf) 15-25% 75-85% AMT1;3 in leaf mesophyll; roots play minor role 5-15 0.1-0.3
PO₄³⁻ PHT1 85-90% 10-15% Limited root uptake; leaves use acid phosphatases & low-affinity P transporters 1-5 0.1-0.5
K⁺ (high) HAK5-L 5-10% 90-95% Leaf-isoform of HAK5; induced under K⁺ deficiency 5-30 0.2-1.2
K⁺ (low) AKT2/3 0% 100% Voltage-gated channels in leaf mesophyll; constitutive expression 50-100 2-4
Fe²⁺ ZIP4 5-10% 90-95% Leaf epidermis expresses high-affinity ZIP4; root IRT1 absent/downregulated 0.5-5.0 0.03-0.3
Fe³⁺ FRO Homologs (Leaf) 0% 100% Plasma membrane ferric reductases in leaf cells (ascorbate-dependent) n/a n/a
Zn²⁺ ZIP7 (Leaf) 5-15% 85-95% ZIP7 expressed in leaf epidermal cells 0.3-2.0 0.02-0.13
Cu²⁺ COPT3 (Leaf) 20-30% 70-80% Chloroplast-targeted COPT3 dominates in leaves 0.05-0.2 0.003-0.013

What Does Km Represent?

You can think of Km as the performance or efficiency of a doorman whose job is to let guests invited to a party into the house.

  • A doorman with a low Km number (in the order of a few µM) is a highly efficient worker who personally attends to each invited guest, but due to this above-standard care, he can only handle a relatively small number of guests.
  • In contrast, a doorman with a high Km number (in the order of hundreds to thousands of µM) is not very efficient, but can handle hundreds to thousands of guests.

In other words, a doorman with a low Km is like a watchmaker with tweezers, while a doorman with a high Km is like a worker with a shovel (for a worker with a shovel, individual components or grains of sand are below his resolution).

Comparing Km values [of high-affinity transporters] between roots and leaves tells you which organ is better equipped to take up a nutrient at very low concentrations.

Critical Notes & Species-Specific Variations

  • Iron:
    • Fe2+ favors root uptake (via IRT transporters).
    • Fe(OH)3 colloids favor leaf uptake (via adhesion/reduction).
    • Fe3+-chelates (e.g., EDTA) split evenly
  • Micronutrients (Zn, Cu): Similar to Fe (root-focused).
  • N/P/K: Root-dominated (75–95%) in all cases due to high-affinity root transporters (NRT, PHT, HAK).
  • Ca2+/Mg2+: Leaf-dominated (60–80%) via diffusion/channels.
  • Morphological Adaptations:
    • Vallisneria: Thick roots → high sediment Fe scavenging (95% root uptake in sediment).
    • Cabomba: Dissected leaves → high foliar efficiency (60% leaf uptake when floating).
    • Myriophyllum: Secretes root/leaf reductants → balanced uptake (50:50 when floating).
  • Exceptions:
    • Rootless species (e.g., Ceratophyllum):
      • Evolved leaf-based Fe3+ reduction (ascorbate-dependent).
      • ZIP transporter density 2× higher than root-dependent species.
    • Oligotrophic specialists (e.g., Lobelia dortmanna): Leaves absorb Fe-colloids via endocytosis.
    • Floating-Leaved Plants (e.g., Nymphaea): Stomatal uptake of Fe-chelate aerosols.
  • Water Chemistry Influence:
    • Low pH (4–5) → ↑ dissolved Fe2+ → ↑ root uptake.
    • High DOC → ↑ Fe3+-OM complexes → ↑ leaf uptake.
  • Sediment-anchored vs Floating roots (no sediment):
    • If the rooting plants do not have access to organic (nutritious) substrate, the ratio of individual transporters between the roots and leaves will further decrease → this means that such plants will be able to obtain fewer nutrients through their roots than if they were rooted in substrate (which is logical). The main reason for this is the limited ability of roots to modify their immediate surroundings (rhizosphere) in flowing water using root exudates.

Why Leaves Struggle with Iron (Compared to Roots):

  1. No Ferric Reductase System: Leaves lack FRO enzymes → cannot reduce Fe3+ efficiently.
  2. Lower Transporter Affinity: Leaf ZIP transporters have Km values 5–10× higher than root IRT1.
  3. Competing Priorities: Leaves prioritize photosynthesis over ion uptake (limited energy allocation).
  4. Oxidative Damage Risk: Leaf Fe2+ uptake triggers ROS (Fenton reactions) without root-like antioxidant systems.

2. Iron Behavior in Slightly Acidic, Oligotrophic, High-Oxygen Natural Waters
(pH ~5, Eh +300 to +400 mV)

  1. Predominant Forms & Speciation:

    • Dominant Oxidation State: Fe(III). The highly oxidizing conditions (Eh +300 to +400 mV) strongly favor Fe3+ over Fe2+. Fe2+ oxidation is rapid under these conditions, especially at pH 5 (kinetics discussed later).
    • Solubility Limitation: Despite the acidic pH, inorganic Fe(III) solubility is extremely low due to hydrolysis and precipitation:
      • Primary Solid Phase: Amorphous Ferrihydrite (Fe(OH)3·nH2O) or freshly precipitated FeOOH. Crystalline forms like hematite or goethite form too slowly to be significant initially.
      • Equilibrium Calculations: Inorganic [Fe3+] at pH 5 is on the order of 10−10 to 10−11 M (0.000006 to 0.00000006 mg/ℓ) – essentially negligible.
    • Crucial Role of Organic Matter (OM): Dissolved Organic Carbon (DOC), primarily humic and fulvic acids, is the key player enabling measurable dissolved Fe.
      • Complexation: OM forms strong soluble complexes with Fe(III). These complexes dominate the truly dissolved (<0.2 µm or <0.45 µm filtered) iron pool.
      • Colloidal Forms: OM also stabilizes Fe(III) (oxyhydr)oxide colloids (nanoparticles). These colloids often pass through filters and contribute to the "operationally defined" dissolved fraction. The boundary between truly dissolved complexes and colloidal Fe-OM is often blurred.
    • Resulting Speciation: The measurable dissolved Fe fraction consists primarily of:
      • Fe(III)-Organic Complexes (e.g., Fe-Humate, Fe-Fulvate)
      • Fe(III)-Organic Colloids
      • Negligible free Fe3+ or Fe2+ ions.
    • Particulate Iron: Consists of freshly precipitated Fe(III) (oxyhydr)oxides, older more crystalline phases, Fe adsorbed onto mineral particles (clays, silt), and Fe associated with particulate organic matter (POM).

  2. Sources (Sites) of Iron:

    • Weathering: Primary source. Chemical weathering of Fe-bearing minerals (e.g., pyrite FeS2, ferromagnesian silicates like olivine/biotite, hematite Fe2O3, magnetite Fe3O4) in catchment soils and bedrock. Acidic conditions enhance silicate weathering. Pyrite weathering releases Fe2+ and sulfate, but rapid oxidation occurs in oxic water.
    • Atmospheric Deposition: Wet (rain, snow) and dry deposition. Can include soluble Fe(II) from combustion or Fe(III) associated with dust particles. Significant source in remote areas.
    • Groundwater Inflow: Often contains elevated Fe(II), especially in anoxic aquifers. Upon entering the oxic water body, rapid oxidation and precipitation occur, potentially forming localized Fe(III) deposits near seepage points.
    • Sediment Diffusion/Resuspension: Underlying sediments are the major reservoir. In *anoxic* sediment porewaters, Fe(II) concentrations can be high. Diffusion of Fe(II) upwards into the oxic water column leads to immediate oxidation/precipitation at the sediment-water interface. Physical resuspension events introduce particulate Fe(III) from sediments.
    • Decomposition: Release of organically complexed Fe during microbial decomposition of plant and algal biomass.

  3. Uptake by Submerged Aquatic Plants:

    • Primary Bioavailable Form: While Fe(III) dominates chemically, submerged plants primarily take up Fe(II). This necessitates a reduction step.
    • Mechanism - Strategy I (Similar to Terrestrial Dicots):
      1. Acidification: H+-ATPases pump protons into the rhizosphere/apoplast, lowering pH locally (can be significant, down to pH 3-4 near roots).
      2. Reduction: Membrane-bound Ferric Chelate Reductase (FCR) enzymes (e.g., FRO family) reduce Fe(III) (whether complexed to OM or inorganic) to Fe(II). This reduction is greatly enhanced by the localized acidification. Reduction kinetics are fast (seconds to minutes).
      3. Transport: Fe(II) is transported across the root plasma membrane via specific Fe(II) transporters (e.g., IRT family - Iron Regulated Transporters).
    • Role of Root Exudates: Plants release phytosiderophores (less common in aquatics than grasses) or, more importantly, low molecular weight organic acids (LMWOAs) (citrate, malate, oxalate) and phenolics. These acids:
      • Chelate Fe(III), improving its solubility and diffusion towards the root.
      • Facilitate reduction by FCR.
      • May directly solubilize Fe from solid phases via ligand-promoted dissolution.
    • Leaf Uptake: Possible, especially for truly dissolved Fe(III)-OM complexes or potentially via foliar reduction, but generally considered less significant than root uptake for submerged macrophytes under typical conditions. May be more important in highly stained (high DOC) waters.
    • Challenge in Oxic Waters: The combination of high O2 (favoring Fe(III)) and low DOC (in oligotrophic systems) makes Fe inherently scarce and poorly soluble. Plants invest significant energy in acidification and reduction to acquire Fe.

3. Adding 0.5 mg/ℓ Fe (as FeCl3 or FeSO4)

A. Addition as FeCl3 (Fe(III)):

  1. Immediate Hydrolysis & Precipitation:
    • Fe3+ undergoes extremely rapid hydrolysis: Fe3+ + H2O ⇌ FeOH2+ + H+; FeOH2+ + H2O ⇌ Fe(OH)2+ + H+; 2FeOH2+ ⇌ Fe2(OH)24+; etc.
    • These monomers and oligomers quickly polymerize into amorphous ferrihydrite particles: Fe3+ + 3H2O → Fe(OH)3 (s) + 3H+.
    • Kinetics: Hydrolysis/polymerization is extremely fast at pH 5. Initial precipitation occurs within seconds to minutes. Aggregation and growth of colloids continue over minutes to hours.
    • Effect on Water Chemistry: The hydrolysis releases H+, causing a localized, transient pH drop near the point of addition. The extent depends on buffering capacity (alkalinity/DOC). In low-buffered oligotrophic waters, a measurable overall pH drop is possible. 0.5 mg/ℓ Fe3+ hydrolysis releases ~0.027 mmol H+/ℓ, which can lower pH significantly in poorly buffered water.
  2. Fate of Precipitated Fe:
    • Forms amorphous ferrihydrite colloids and particles.
    • Particles aggregate (coagulation) over time (minutes to hours), forming larger flocs that settle. Settling rate depends on particle size, turbulence, and density.
    • Interaction with Organic Matter: Dissolved OM will rapidly adsorb onto the fresh ferrihydrite surfaces or form complexes with soluble hydrolysis products before precipitation. This OM coating:
      • Slows down further aggregation and settling.
      • Stabilizes smaller colloidal particles.
      • May slightly increase the dissolved (<0.45 µm) Fe fraction measured shortly after addition compared to OM-free water.
  3. Resulting Forms (Over Time):
    • Seconds: Hydrolyzed monomers/oligomers, initial nanocolloids.
    • Minutes: Amorphous ferrihydrite colloids (OM-coated), aggregating flocs.
    • Hours-Days: Settling flocs, adsorbed Fe onto existing suspended particles, Fe incorporated into biofilms. A small fraction of Fe(III)-OM complexes may remain in solution. Very little truly dissolved inorganic Fe.

B. Addition as FeSO4 (Fe(II)):

  1. Immediate Oxidation:
    • Fe2+ is rapidly oxidized by dissolved oxygen: 4Fe2+ + O2 + 4H+ → 4Fe3+ + 2H2O
    • Kinetics: Oxidation rate of Fe2+ is highly pH-dependent. At pH 5, it is relatively slow compared to neutral/alkaline pH. The half-life (t₁/2) of Fe2+ can range from minutes to several hours at pH 5 and 20°C, depending on exact conditions (ionic strength, catalysis). Oxidation is autocatalytic (Fe(III) products catalyze further oxidation). Presence of OM or surfaces can alter kinetics.
  2. Hydrolysis & Precipitation:
    • As Fe2+ oxidizes to Fe3+, the newly formed Fe3+ immediately undergoes hydrolysis and precipitation as described for FeCl3 addition. The precipitation kinetics are similarly fast.
  3. Transient Effects:
    • pH: Oxidation consumes H+, causing a localized, transient pH increase near the point of addition. 0.5 mg/ℓ Fe2+ oxidation consumes ~0.009 mmol H+/ℓ. This pH rise is generally smaller than the drop caused by FeCl3 hydrolysis.
    • Dissolved Fe(II): A measurable pool of dissolved Fe(II) will persist for minutes to hours before complete oxidation.
  4. Resulting Forms (Over Time):
    • Seconds-Minutes: Dissolved Fe(II), initial Fe(III) hydrolysis products from oxidized Fe.
    • Minutes-Hours: Decreasing dissolved Fe(II), increasing amorphous ferrihydrite colloids/flocs (OM-coated).
    • Hours-Days: Identical to FeCl3 addition: Settling flocs, adsorbed Fe, minor Fe(III)-OM complexes. Essentially indistinguishable from the FeCl3 endpoint.

Summary of Addition:

Regardless of initial form (FeCl3 or FeSO4), within hours to a day in oxic, pH ~5 water, the vast majority (>95-99%) of the added 0.5 mg/ℓ Fe will be present as amorphous ferrihydrite particles/colloids, either settling or suspended. A small fraction will be stabilized as Fe(III)-OM colloids/complexes or adsorbed onto existing particles. Very little will remain as truly dissolved Fe(II) or Fe(III) ions. The measurable "dissolved" Fe fraction (<0.45 µm) shortly after addition will consist mainly of Fe(III)-OM complexes and stabilized colloids, but this will decrease as larger aggregates form and settle. The transient pH effects (drop for FeCl3, rise for FeSO4) are noticeable in poorly buffered waters.

4. Toxicity and Reactive Oxygen Species (ROS)

  1. Toxicity of Iron Forms:
    • Fe(II) Toxicity: This is the primary concern for direct phytotoxicity in aquatic plants.
      • Mechanism 1: Fenton Reaction: Fe2+ + H2O2 → Fe3+ + •OH + OH− (Hydroxyl Radical). •OH is an extremely potent ROS, causing indiscriminate damage to lipids (membranes), proteins, and DNA.
      • Mechanism 2: Haber-Weiss Reaction: Fe3+ + O2•− → Fe2+ + O2 followed by Fenton. Amplifies ROS production using superoxide.
      • Consequences: Lipid peroxidation, enzyme inactivation, chlorophyll degradation, membrane leakage, growth inhibition, cell death. Symptoms include bronzing/reddening, necrotic spots, root browning, reduced growth.
      • Threshold: Toxicity thresholds vary greatly by species and conditions. Levels as low as 0.5-2 mg/ℓ Fe2+ can cause toxicity in sensitive species under prolonged exposure, though higher levels are often tolerated acutely. The addition of 0.5 mg/ℓ total Fe would initially generate significant dissolved Fe(II) only if added as FeSO4, and its persistence depends on oxidation kinetics.
    • Fe(III) Toxicity: Inorganic Fe(III) (free ions, colloids, particles) has very low direct phytotoxicity due to poor bioavailability and inability to directly drive Fenton reactions. However:
      • Indirect Physical Effects: High loads of Fe(III) precipitates can:
        • Smother plants/roots/biofilms, impairing gas exchange and light penetration.
        • Clog stomata (on floating leaves) or cuticles.
        • Adsorb essential nutrients (e.g., P), inducing deficiencies.
      • Potential for ROS: If Fe(III) is reduced intracellularly or at the root surface without adequate antioxidant capacity, it can generate Fe(II) and subsequently ROS. This is more likely a consequence of uptake dysregulation under excess than direct toxicity of external Fe(III).
  2. ROS Formation Conditions:
    • Essential Requirements: Simultaneous presence of Fe(II) (or reducible Fe(III)) and H2O2/O2•− (Reactive Oxygen Species precursors).
    • Key Scenarios in Plants:
      • Excess External Fe(II): Directly fuels extracellular Fenton/Haber-Weiss reactions. Most relevant shortly after FeSO4 addition.
      • Internal Fe Overload: If uptake/reduction exceeds cellular demand and sequestration capacity (ferritin), free Fe(II) can accumulate in the cytosol/chloroplasts. Reaction with endogenous H2O2/O2•− (produced normally during metabolism or under stress) leads to internal oxidative damage.
        • This scenario could be the most likely cause of growth deformities [due to iron poisoning] in some more sensitive plant species such as R. wallichii or A. pedicellata ‘Gold’.
      • Reductive Stress: Under conditions causing over-reduction of cellular components (e.g., low light, certain pollutants), excessive reduction of internal Fe(III) stores or external Fe(III) at the root surface can generate Fe(II) and ROS.
      • Apoplastic Space: The root apoplast (cell wall space) is a potential site for Fe(II)/H2O2 interaction, especially if Fe uptake/reduction is dysregulated.
    • Defense Mechanisms: Plants deploy antioxidants (SOD, CAT, APX, GPX, Ascorbate, Glutathione) and Fe-sequestration proteins (Ferritin) to mitigate ROS and buffer Fe.

Relevant References & Keywords for Further Research:

  • Speciation/Kinetics: Stumm & Morgan (Aquatic Chemistry), Davison & De Vitre (Iron redox kinetics), Ferrihydrite, NOM-Fe complexation, Fe(II) oxidation kinetics (pH dependence).
  • Plant Uptake: Strategy I, Ferric Chelate Reductase (FRO), IRT transporters, Root exudates (LMWOAs), Iron plaque.
  • Toxicity/ROS: Fenton reaction, Haber-Weiss cycle, Lipid peroxidation (MDA assay), Antioxidant enzymes (SOD, CAT, APX, GPX), Ferritin.
  • Aquatic Plant Focus: Look for work by researchers like J. E. Vermaat, W. Armstrong, S. Maberly, or specific studies on Lobelia dortmanna, Isoetes spp., or Eriocaulon spp. which often inhabit oligotrophic, acidic waters. Search terms: "Iron acquisition submerged macrophytes", "Iron toxicity aquatic plants", "ROS aquatic macrophytes iron".

Tricky microelements

Microelements toxicity despite their chelation

There seems to be an apparent contradiction: While microelements should be more available (and therefore potentially more toxic) at low pH, they are actually most toxic at higher pH, even when chelated. How is this genuine paradox in plant nutrition (i.e. the disconnect between theoretical availability and actual uptake) possible?

Chelates breakdown under higher pH & bicarbonate presence

The key here is understanding that "availability" isn't just about solubility. At higher pH, while free Fe³⁺/Mn²⁺ decrease, the chelates themselves behave differently. DTPA's stability plummets above pH 6.5, and bicarbonate actively attacks metal-chelate bonds. I picture bicarbonate ions swarming the Fe-DTPA complexes like little demolition crews breaking them apart.

Precipitation pathway

When chelates fail, the resulting ferric hydroxide isn't just inert sediment. Under oxygenated conditions, it forms reactive nanoparticles that stick to roots and leaves – like microscopic shards of rust that plants accidentally ingest. Our oxygen-saturated aquariums actually worsens this! More oxygen means faster oxidation of any released Fe²⁺ to these problematic Fe³⁺ colloids. Note the oxygen paradox: our well-aerated tanks, intended to help plants, actually create the perfect storm for Fe/Mn toxicity via oxidation and precipitation.

Manganese escape and confusion with magnesium

For manganese, the story is different but equally treacherous. Mn-EDTA is inherently weaker [than Fe-DTPA], and in bicarbonate-rich tanks, it's like the Mn ions are constantly escaping custody. Plants then overabsorb Mn because it's chemically similar to magnesium they desperately need. I imagine confused leaf/root cells grabbing "imposter" manganese instead of precious magnesium.

Organic substrates contribution

And the organic substrates? They're sneaky contributors too – decomposing organics release compounds that further destabilize chelates.

The resulting toxicity

The result is not just toxicity; it's a plant choking on rust particles. So here's the whole path: failed chelation → nanoparticle formation → forced uptake → oxidative havoc inside plants. Only in acidic tanks do the microelements avoid this fate. So now you may understand why our scientific intuition ("high pH should reduce availability") was tricked by these complex chemical behaviors.

More detailed explanation

Chelate instability & bicarbonate interference at higher pH

  • The Problem: Fe-DTPA and Mn-EDTA are strong chelates, but their stability is pH-dependent.

    • Fe-DTPA: Optimal stability is between pH 3.0 and 6.0. Above pH 6.0, stability decreases significantly. Above pH 6.5-7.0, DTPA struggles to hold Fe³⁺ effectively.
    • Mn-EDTA: Stability is generally lower than Fe-DTPA and also decreases above neutral pH. EDTA is a weaker chelator for Mn²⁺ than DTPA is for Fe³⁺.

  • Bicarbonate's crucial role: Bicarbonate (HCO₃⁻) is a potent chelator antagonist:

    • It competes with the chelate for binding sites on Fe/Mn ions.
    • It accelerates the hydrolysis and oxidation of Fe³⁺ released from unstable chelates.
    • It directly increases the pH locally around roots or surfaces, further destabilizing chelates.

  • Result: In tanks with pH ≥6 and bicarbonate, the Fe-DTPA and Mn-EDTA complexes became unstable. Fe³⁺ and Mn²⁺ ions are then partially displaced or released, making them vulnerable to precipitation.

Formation of highly reactive and bioavailable precipitates

  • What precipitates form?

    • Iron: Released Fe³⁺ rapidly hydrolyzes and oxidizes (especially in oxygenated water) forming amorphous ferric (oxy)hydroxides - Fe(OH)₃, FeOOH. These initially form as nanoparticles or colloids.
    • Manganese: Mn²⁺ oxidizes more slowly than Fe²⁺ but can form Mn(III/IV) oxides/hydroxides (MnOOH, MnO₂) in oxygenated conditions, especially on surfaces or catalyzed by bacteria.

  • Why "highly reactive"?

    • These freshly precipitated nanoparticles and colloids have an immense surface area-to-volume ratio.
    • Their surfaces are covered with reactive hydroxyl groups (-OH) and positive charges (especially Fe oxides at near-neutral pH), making them highly adsorbent.
    • They are thermodynamically unstable intermediates striving to form larger, more stable crystals (like rust), releasing energy in the process.

  • Why "bioavailable"? (The toxicity paradox)

    1. Direct Root Contact & Adsorption: Plant roots constantly release organic acids, protons (H⁺), and reductants to solubilize nutrients. These exudates aggressively dissolve freshly precipitated Fe/Mn oxides/nanoparticles adhered to root surfaces or rhizosphere particles. The plant absorbs the dissolved Fe²⁺/Mn²⁺ faster than it can regulate.
    2. Colloidal Uptake: Extremely small nanoparticles/colloids can be directly taken up by roots via endocytosis or through cracks in the root epidermis/apoplast.
    3. Chemical Reduction: Root exudates chemically reduce Fe³⁺ in oxides to soluble Fe²⁺ right at the root surface, facilitating uptake. Plants have dedicated ferric reductases for this purpose, but they can be overwhelmed by a massive local supply.
    4. Bypassing Regulation: Plants tightly regulate uptake of soluble Fe²⁺/Mn²⁺ ions via transporters. However, the local dissolution of massive amounts of precipitate at the root surface creates a concentrated "pulse" of ions that can flood into the root faster than regulatory mechanisms can respond. It's akin to force-feeding.
    5. Oxidative Stress: Once inside the plant, excess Fe/Mn catalyze the production of highly damaging reactive oxygen species (ROS) via Fenton/Haber-Weiss reactions (Fe²⁺ + H₂O₂ -> Fe³⁺ + OH• + OH⁻). This internal oxidative damage is a primary cause of toxicity symptoms (chlorosis, necrosis, stunting).

  • Oxygenation Paradox: Highly oxygenated water in planted tanks exacerbates this issue.

    • It rapidly oxidizes any Fe²⁺ released from unstable chelates or by root reduction, driving the formation of the reactive Fe(III) precipitates/colloids near the roots.
    • While O₂ is good for root respiration, it fuels the precipitation process that leads to the bioavailable forms causing toxicity.

Contrast with low pH conditions

  • Chelate Stability: At pH 4.5, Fe-DTPA and Mn-EDTA are exceptionally stable. DTPA holds Fe³⁺ tightly, and EDTA holds Mn²⁺ effectively.
  • Solubility: Fe³⁺ and Mn²⁺ remain soluble at low pH. Precipitation of oxides/hydroxides does not occur.
  • Controlled Release: The stable chelates act as true slow-release reservoirs. Plants can efficiently take up Fe/Mn via regulated reduction (Fe) or transport (Mn) mechanisms only as needed. The chelate prevents a flood of ions.
  • Result: Tissue concentrations remain within sufficient ranges (Fe ~0.01-0.05%, Mn ~0.005-0.05%), avoiding deficiency and toxicity. This explains the excellent plant performance in tanks with acidic water despite the low pH.

Practical recommendations

How to maintain micronutrient availability without triggering toxicity when chemistry works against us?

  • 1. Iron solution: Upgrade to Fe-EDDHA

    • Why it works:
      • EDDHA maintains iron solubility up to pH 9-11 due to its ortho-ortho' hydroxy group configuration, resisting bicarbonate interference.
      • Note: EDDHA may tint water slightly red; this is harmless.

  • 2. Manganese challenge: Why EDTA fails & alternatives

    • With manganese (a trickier element to chelate) EDTA really is the standard, but its pH weakness is the problem.
    • This problem can be approached from three angles: better chelates (if any exist), dosage/timing tweaks, and environmental hacks.
      • Mn-EDTA stability collapses above pH 6.5. Bicarbonate displaces Mn²⁺, leading to oxidation → MnO₂ precipitation → root-zone toxicity.
      • Recommendation: Use Mn-EDDHSA (if available)
        • EDDHA/EDDHSA resist bicarbonate and high pH.
      • If Mn-EDDHSA unavailable:
        • Using DTPA (over EDTA) could improve Mn stability, but only marginally (not dramatically).
        • Lowering phosphate might help reduce precipitation.
        • Using humic acids [as natural chelators] may help keep the manganese in solution while not toxic.
          • Combine Mn with humic/fulvic acids (20-50 ppm) – they complex Mn and stabilize colloids.

Toxicity mechanisms despite low external concentrations

Cellular iron overload:

  • Fenton chemistry: Excess Fe²⁺ catalyzes hydroxyl radical formation:
    • Fe²⁺ + H₂O₂ → Fe³⁺ + OH• + OH⁻
  • Organelle damage: Fe accumulates in chloroplasts → lipid peroxidation → photosystem degradation
  • Protein oxidation: Essential enzymes (RuBisCO) become non-functional

Manganese toxicity:

  • Competitive inhibition: Excess Mn²⁺ displaces Mg²⁺ in photosystem II
  • Oxidative stress: Mn accumulation in leaf margins → necrotic spots

Bicarbonate effects:

  • Competitive binding: Bicarbonate competes with synthetic chelates for metal coordination
    • Bicarbonate actively attacks metal-chelate bonds. Bicarbonate ions are swarming the chelate complexes like little demolition crews breaking them apart. When chelates fail, the resulting hydroxides form reactive nanoparticles under oxygenated conditions that may cause oxidative havoc inside plants.
    • Besides this, bicarbonate forms soluble complexes with Mn²⁺ (MnHCO₃⁺) that are more readily absorbed than free Mn²⁺ ions. Thus, while bicarbonate decreases total Mn solubility in bulk solution, it increases Mn bioavailability to plants through complexation.
    • Also high bicarbonate concentrations may cause severe potassium leakage from plant cells
  • pH effects: Bicarbonate buffering can shift pH away from optimal chelate stability ranges
  • Complex formation: Formation of metal-bicarbonate complexes reduces chelated metal availability

Organic substrate effects:

  • Enhanced metal mobilization: Organic acids increase metal solubility and bioavailability + organic matter content releases substantial Fe/Mn during decomposition

Chemistry of iron in natural water

See: Hem & Cropper. Chemistry of iron in natural water. Washington, U.S., 1962 (268 p.)

Note: This diagram is practically useless for aquarists because it only shows the thermodynamic equilibrium state (i.e., the most stable form of iron assuming infinite time) while ignoring kinetics. However, kinetics override the equilibrium state, so the predicted thermodynamic state does not apply here [temporarily]. After all kinetic processes (e.g., oxidation and precipitation) have taken place, the form of iron predicted by the Pourbaix diagram will eventually establish itself in the water. By that time, however, the vast majority of iron will have precipitated, so only a fraction of it will remain dissolved in the water (but this fraction will be in the form predicted by the Pourbaix diagram).

  • amounts of iron in solution [= in dissolved state] in natural water at equilibrium are related to the pH and Eh (= redox potential) of the solution
  • iron occurs in two oxidation states → the divalent or ferrous form (Fe2+) and the trivalent or ferric form (Fe3+)
    • stability-field diagram shows the Eh and pH values at which each of these predominates
      Stability-field diagram for aqueous ferric-ferrous system (ferrous-ferric chemical equilibria and redox potentials).
    • but [in natural waters] the ferrous iron (Fe2+) will be quickly lost from solution by oxidation and precipitation of ferric hydroxide [Fe(OH)3] at a rate governed by the diffusion of oxygen through the water
    • thus, main ionic species present [in natural waters] include Fe3+ only in the form of Fe(OH)3 and/or FeOH2+
    • the amounts of iron that theoretically could be present in solution are mostly bellow 0.01 ppm [under pH 5-8 and Eh 300-500 mV]
      • in aerated waters whose pH is above ~5, ferric iron (Fe3+) can be present in excess of 0.01 ppm only as a suspension of oxide or hydroxide [or in organic complexes]
  • iron in aqueous solution is subject to hydrolysis
  • the iron hydroxides [Fe(OH)n] formed in these reactions (especially the ferric form = Fe(OH)3) have very low solubility
    • at equilibrium [in the pH range of 5 to 8] this compound is largely in the solid state
  • in most natural waters, the pH is not low enough to prevent hydroxides from forming, and under oxidizing conditions, practically all the iron is precipitated as Fe(OH)3 (or more correcly, Fe2O3*3H2O)
  • ferric iron is a powerful former of complexes [with inorganic as well as organic material]
    • these ions may be considerably more stable than the uncomplexed iron and more may remain in solution
    • inorganic material (ions) with which ferric iron forms complexes: chloride, fluoride, phosphate, sulfate, and carbonate
    • organic material: humic acids

Iron Species Kinetics in Soft Water

Under normal conditions (= pH > 5.0 & redox around +400 mV) iron in water will quickly oxidize and precipitate – as Fe(OH)3/FeOOH – unless chelated.

Kinetic persistence of iron species in soft water (pH 5-6, Eh +200 to +500 mV)

Practical implications:

  • In soft, oxygenated water at pH 5–6 and high redox (i.e., typical parameters of planted aquariums), Fe²⁺ disappears within minutes to an hour, replaced by insoluble Fe(OH)₃(s).
    • at pH 4.0 & Eh +400 mV
      • dominant form: Fe²⁺(aq) (>99.9%)
      • oxidation: very slow (half-life: days-weeks)
      • precipitation: no
    • at pH 5.0 & Eh +400 mV
      • dominant form: Fe²⁺(aq) (50-80%)
      • oxidation: significant (half-life: tens of minutes)
      • precipitation: rapid
    • at pH 5.5 & Eh +400 mV
      • dominant form: Fe(OH)₃/FeOOH (>95%)
      • oxidation: very fast (half-life: minutes)
      • precipitation: near-complete
    • at pH 6.0 & Eh +400 mV
      • dominant form: Fe(OH)₃/FeOOH (>99%)
      • oxidation: extremely fast (half-life: <minutes)
      • precipitation: immediate
  • Thermodynamic diagrams (like Pourbaix) show what is stable at equilibrium, but in real water, iron's kinetic persistence is brief—soluble iron is quickly removed from solution by oxidation and precipitation.

At pH 5 [in 15 minutes] only about ⅒ of the iron is said to be oxidized, at pH 6 then ½ of the iron, and at pH 7 and higher practically all the iron. Although I have not been able to verify this information, if it is true, it gives us some idea of how long unchelated iron is normally available to plants in the aquarium at different pH.

At higher pH and redox, unchelated iron in water oxidises and precipitates more quickly, so plants have less time to absorb it, which means they can suffer much more from iron deficiency at higher pH/redox.

Iron in Substrate

First, note the difference between terrestrial substrates and aquatic sediments:

  • terrestrial substrates
    • have a high (positive) redox [usually from +400 to +700 mV]
      • most nutrients in them therefore remain in an oxidized (= inaccessible) state
      • terrestrial plants must therefore reduce (and thus make available) the oxidised nutrients in the substrate themselves by releasing their own reducing substances (= special organic acids)
  • aquatic sediments
    • have a low (negative) redox [usually from −250 to −300 mV]
      • most of the oxidised nutrients are therefore naturally reduced (= dissolved and available to plants)
      • aquatic plants thus have nutrients naturally available in abundance in aquatic sediments and do not have to actively worry about making them available; rather, they have to actively limit their excessive availability [to a safe level] (= not to be poisoned)

Thus, a completely opposite situation (than in water) occurs in the aquatic sediment. Here, under normal conditions, within a few weeks of flooding, a dramatic drop in redox occurs and a strongly reducing environment is created in which most of the oxidised compounds (including iron) begin to dissolve and thus become available to the plants.

So, over time, the problem here is not with iron (and other micronutrients) deficiency, but with its toxicity. There are so many dissolved microelements in the soil solution that the plants are at risk of poisoning. Plants deal with this danger of poisoning in a quite simple way → by releasing oxygen into the immediate surroundings of their roots (the so-called rhizosphere), whereby they partially oxidize the excess of reduced (= accessible) microelements again, and thus practically make them inaccessible again. Some plants also use other substances (e.g. various organic compounds) to neutralise the dissolved microelements.

And while we can do something about the impending iron deficiency in the water (e.g. use some chelate), we can't do much about the impending heavy metal poisoning in the substrate. The only thing we can do is (1) choose a substrate that contains a normal (reasonable) amount of nutrients and (2) ensure that the plants have well oxygenated water available at all times from which they can pump oxygen into the rhizosphere.

See: Aldridge & Ganf. Modification of sediment redox potential by three contrasting macrophytes. Marine and Freshwater Research. 2003, 54(1), 87-94. ISSN 1323-1650. DOI 10.1071/MF02087.

Change in sediment redox potential beneath Potamogeton crispus over time since inundation measured at depth of 2 cm. After about 14 days, the redox potential of the flooded sediment dropped to −150 mV and continued to decline over time. Thus, even at this shallow depth, the redox in the sediment reaches significantly negative values at which most of the oxidized nutrients will be reduced (i.e. dissolved and made available to plants).
Redox potential measured beneath different aquatic plants and bare sediment (= sediment without plants) at depth of 1 and 10 cm. The first two plants (Typha domingensis and Bolboschoenus caldwellii) grew above the water level; Potamogeton crispus was growing fully submerged. In the case of T. domingensis it had a significant effect on the oxygenation of the sediment and an increase in its redox potential. For fully submerged aquatic plants (including aquarium plants), the redox shows significantly negative values.

See: Ferreira & Knupp. Chemistry of Lowland Rice Soils and Nutrient Availability. Communications in soil science and plant analysis. 2011, 42(16), 1913-1933. ISSN 0010-3624 print / 1532-2416 online. DOI 10.1080/00103624.2011.591467.

When the soil is flooded, it quickly develops two diametrically opposed layers: On the top is an oxidized or aerobic surface layer where oxygen is present, with a reduced or anaerobic layer underneath in which no free oxygen is present. Illustrated in the picture below is the thin oxidized layer (usually 1 to 20 mm in thickness), beneath which is a thick reduced layer of soil. In addition, flooding also has major effects on the availability of macro and micronutrients. Some nutrients are increased in availability to the plants, whereas others are subject to greater fixation or loss from the soil as a result of flooding.

Within 6-10 h after flooding, the O2 level drops to near zero. The rapid declines of O2 from the soil are accompanied by an increase of other gases produced through the microbial respiration. The major gases that accumulate in the flooded soils are carbon dioxide (CO2), methane (CH4), nitrogen (N2), and hydrogen (H2). The composition of these gases may vary from 1 to 20% CO2, 10 to 95% N2, 15 to 75% CH4, and 0 to 10% H2. This variation may be associated with the presence of microbial biomass, organic matter, and inorganic substances and also the plant species planted. In flooded soils, aerobic microorganisms become quiescent or die, and facultative and obligate anaerobic bacteria proliferate. These new microorganisms oxidize organic compounds with the release of energy in a process called “anaerobic fermentation”.

The main changes that occur in flooded soils are decreases in redox potential and increases in iron (Fe2+) and manganese (Mn2+) concentrations because of the reductions of Fe3+ to Fe2+ and Mn4+ to Mn2+.

  • Fe2O3 + 6H+ + 2e− ↔ 2Fe2+ + 3H2O
  • MnO2 + 4H+ + 2e− ↔ Mn2+ + 2H2O

The pH of acidic soils increases and alkaline soils decreases [by up to about 1 unit] as a result of flooding. Overall, pH of most soils tends to change toward neutral after flooding.

The change in soil pH of acidic soils collected from four locations in Rio Grande, Brazil. As can be seen, the soil pH increases with flooding and stabilizes around 56 days after flooding in all the soils.

Other results are the reduction of nitrate (NO3−) and nitrogen dioxide (NO2−) to dinitrogen (N2) and nitrous oxide (N2O); reduction of sulfate (SO42−) to sulfide (S2−); reduction of carbon dioxide (CO2) to methane (CH4); improvement in the concentration and availability of phosphorus (P), calcium (Ca), magnesium (Mg), iron (Fe), manganese (Mn), molybdenum (Mo), and silicon (Si); and decrease in concentration and availability of zinc (Zn), copper (Cu), and sulfur (S).

Reduction processes taking place in submerged soils:

ReactionRedox
(mV)
O2 + 4H+ + 4e− ↔ 2H2O+810
2NO3− + 12H+ + 10e− ↔ N2 + 6H2O+740
MnO2 + 4H+ + 2e− ↔ Mn2+ + 2H2O+400
CH3COCOOH + 2H+ + 2e− ↔ CH3CHOHCOOH−160
Fe(OH)3 + 3H+ + e− ↔ Fe2+ + 3H2O−190
SO42− + 10H+ + 8e− ↔ H2S + 4H2O−210
CO2 + 8H+ + 8e− ↔ CH4 + 2H2O−240
N2 + 8H+ + 6e− ↔ 2NH4+−280
NADP+ + 2H+ + 2e− ↔ NADPH−320
NAD+ + 2H+ + 2e− ↔ NADH−330
2H+ + 2e− ↔ H2−410
Ferredoxin (ox) + e− ↔ Ferrodoxin (red)−430

Note on toxic hydrogen sulfide (H2S) production: In flooded soils, SO42− ion is reduced to toxic hydrogen sulfide (H2S) by anaerobic microbial activities. However, since by the time this starts to happen, reduced iron (Fe2+) will already be present in the soil [the reduction of which always precedes the reduction of sulphates], the H2S produced will be converted to insoluble iron sulfide (FeS). Thus, this reaction protects microorganisms and higher plants from the toxic effects of H2S. Due to this, the availability of S is reduced in flooded soils due to formation of insoluble FeS.

Influence of flooding on redox potential of some Mexican soils.

After flooding, the soil also leaches out the released nutrients, which increases its conductivity (see figure below).

Influence of flooding on electrical conductivity of some Mexican soils.
PS: 1 millimho (mmho) = 1 milliSiemens (mS) = 1,000 microSiemens (µS)

As far as the total iron content in soil is concerned, the lower limit (when there is usually a risk of iron deficiency) is 2.5-5.0 mg/kg Fe, and the upper limit (which can be toxic for some plants under certain conditions) is 300 mg/kg.

Iron Concentration in Aquatic Plants

For an idea of what iron (Fe) concentrations can be expected [under various conditions] in aquatic macrophytes, see the following overview:

Summary of General Iron Concentrations and Thresholds in Aquatic Plants

CategoryConcentration
Range
(ppm DW)		Sources
Typical/Normal Range100 - 200^1
General Plant Leaf Tissue50 - 250^2
Micronutrient Range0.1 - 200^3
Deficiency Threshold< 50^2
Toxicity Threshold> 500-700^4

Phragmites australis

External
Fe Supply
(mg/ℓ)		Root Mean
(Range)
(ppm)		Rhizome Mean
(Range)
(ppm)		Shoot Mean
(Range)
(ppm)
0719 (50 - 1750)88.6 (BDL - 139)21.9 (BDL - 31.8)
0.05568 (155 - 926)24.6 (4.25 - 67.8)36.1 (10.2 - 86.8)
0.11327 (244 - 2174)151 (21.6 - 279)65 (20.9 - 114)
0.55465 (3428 - 8757)232 (173 - 442)42 (10.4 - 67.7)
1.014608 (8851 - 18070)642 (430 - 846)125 (17.5 - 316)
2.051231 (46734 - 61991)4197 (1125 - 7497)97 (22.1 - 240)
5.041346 (30482 - 56176)6940 (921 - 10415)144 (39 - 253)
10.059096 (42582 - 100930)13174 (8612 - 24691)216 (94 - 399)
20.051072 (52202 - 55691)21818 (17170 - 26582)477 (42 - 871)
50.076913 (50617 - 125520)51846 (16310 - 159174)729 (286 - 1501)



Ceratophyllum demersum

SeasonMean Iron
Concentration
(ppm DW)		Sources
Spring 20052350^1
Summer 20055150^1
Autumn 20051500^1
Winter 20052700^1
Spring 20061400^1
Overall Mean2620^1
Max. in Wastewater4291^2



Potamogeton pectinatus (syn. Stuckenia pectinata)

SeasonMean Iron
Concentration
(ppm DW)		Sources
Spring 20051100^1
Summer 20052280^1
Autumn 2005550^1
Winter 20052250^1
Spring 2006950^1
Overall Mean1430^1
Organ-Specific:
Leaf Tissue1184^2
Root Tissue7166^2



Ottelia alismoides

Plant
Organ		Mean Iron
Concentration
(ppm DW)		Sources
Leaves1340^1
Stem1021^1
Roots1118^1



Elodea/Egeria densa

  • Critical concentration in dry weight: 60.0 ppm Fe
    • This concentration is associated with yields approximately 5% below the maximum, indicating a threshold for optimal growth.
    • Source: nepis.epa.gov



Iron-Phosphorus (Fe-P) Interactions

Let's say we add mg/ℓ Fe (as FeCl₃ or FeSO₄) to the water with mg/ℓ initial H₂PO₄⁻ concentration. How much phosphate (P) would adsorb onto the resulting ferrihydrite (Fe(OH)₃) precipitates? Here's a breakdown of the process and estimation:

Conditions

  • pH ~5
  • Eh +300-400 mV
  • Negligible Alkalinity/DOC

Key Factors Influencing P Adsorption

  • Dominant Iron Phase: Freshly precipitated amorphous ferrihydrite (Fe(OH)₃) forms rapidly. This phase has a very high specific surface area (100-700 m²/g) and reactivity, making it an excellent P adsorbent.
  • pH Dependence: P adsorption onto Fe oxides is strongly pH-dependent.
    • Maximum adsorption occurs between pH 3 and 5.5.
    • At pH ~5, adsorption is near its peak. Below pH 3, H₃PO₄⁰ dominates and adsorbs weakly; above pH 5.5, increasing OH⁻ competition and negative surface charge reduce adsorption.
  • Adsorption Capacity: Ferrihydrite has a high theoretical maximum P adsorption capacity.
    • Reported values range from 0.5 to 1.5 mg P per mg Fe, depending on ferrihydrite structure, aging, and experimental conditions.
    • 0.5 mg P / mg Fe is a widely accepted and conservative estimate for fresh ferrihydrite at pH ~5.
  • Phosphate Availability (Critical Factor):
    • Your system is oligotrophic, meaning dissolved inorganic phosphate (DIP, primarily H₂PO₄⁻ at pH 5) concentrations are *very low*.
    • Typical oligotrophic DIP concentrations range from <0.001 mg/ℓ P to 0.02 mg/ℓ P (1 to 20 µg P/L).
    • This is likely the *limiting factor* for total P adsorption.
  • Competing Anions: Negligible bicarbonate and DOC mean no significant competition for adsorption sites from CO₃²⁻ or humic/fulvic acids.
  • Kinetics: Adsorption of P onto freshly precipitated ferrihydrite is very rapid. Significant adsorption occurs within minutes, approaching equilibrium within hours.

Estimation of P Adsorbed

  • Initial H₂PO₄⁻ concentration = mg/ℓ
    • Convertion to P: 0.32 × mg/ℓ = mg/ℓ P
  • Added Fe = mg/ℓ
  • Ferrihydrite (Fe(OH)₃) adsorption capacity = mg P adsorbed per mg Fe
  • Max P adsorption capacity = mg/ℓ Fe × mg P per mg Fe = mg/ℓ P ÷ 0.32 = mg/ℓ H₂PO₄⁻
    • Formula: H₂PO₄⁻ adsorbed (mg/ℓ) = Fe (mg/ℓ) × (0.5 ÷ 0.32) = Fe (mg/ℓ) × 1.56
    • Result: After adding mg/ℓ Fe to the water containing mg/ℓ H₂PO₄⁻ (i.e., mg/ℓ P), mg/ℓ H₂PO₄⁻ (i.e., mg/ℓ P) is adsorbed onto ferrihydrite (Fe(OH)₃) precipitate, leaving mg/ℓ H₂PO₄⁻ (i.e., mg/ℓ P) in the water.

Key Results Summary Table

ParameterP
(mg/ℓ)		H₂PO₄⁻
(mg/ℓ)		Calculation
Initial DIP (as P)(32% of H₂PO₄⁻ mass)
Fe(OH)₃ adsorption capacity( mg/ℓ Fe × 0.5 mg P per mg Fe)
P Adsorbed
P Remaining
Adsorption Efficiency~%
DIP = dissolved inorganic phosphorus

Fate of Adsorbed P

  • Adsorbed P is sequestered in Fe(OH)₃ particles (Fe–P complexes).
  • Particles settle to sediments, transferring P from the water column.
  • In oxic sediments, P remains tightly bound; in anoxic zones, microbial Fe(III) reduction may release P slowly.
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