Just as there are six macronutrients, so are there six micronutrients: iron (Fe), manganese (Mn), boron (B), zinc (Zn), copper (Cu), and moybdenum (Mo).
Micronutrients are so named as they are needed in smaller amounts than the macronutrients. As they are elements, required in trace amounts, they are also referred to as trace elements.
Iron, the micronutrient we will cover in this post, is the fourth most abundant element in the Earth’s crust after oxygen (O), silicon (Si) and aluminium (Al). It makes up about 5% of the crust by weight, with the macronutrient calcium the fifth most abundant at 3.64%.
Iron in soil comes from weathered rocks, and is also a major component of clay minerals. Iron may be abundant in soil, but the amount that is soluble and therefore available to plants is small in comparison to the total iron content of the soil.
Iron chemistry gets complicated quickly, as iron has two oxidation states. Very simply this means that iron can form ferrous ions (Fe2+) and ferric ions (Fe3+).
In the presence of oxygen, Fe2+ oxidises to Fe3+. The Fe2+ ion loses an electron to become an Fe3+ ion, and its oxidation number increases from 2 to 3:
Fe2+ → Fe3+ + e-
In the absence of oxygen, Fe3+ reduces to Fe2+. The Fe3+ ion gains an electron to become an Fe2+ ion, and its oxidation number decreases from 3 to 2:
Fe3+ + e- → Fe2+
Fe2+ is said to be reduced and Fe3+ is said to be oxidised.
But why am I bothering you with all this? Because both soil chemistry and biochemistry play a major role in determining which iron ions and compounds are present in the soil, and hence ultimately available to plants, as we shall see.
Soluble iron in soil exists as ferrous ions (Fe2+), ferric ions (Fe3+), and the ferric iron hydroxide ions Fe(OH)2+ and FeOH2+.
(It is more correct these days to indicate the oxidation state with Roman numerals. Thus for the above we’d write Fe(II), Fe(III) and iron (III) hydroxide ions respectively.)
Fe2+ ions are not as common in well-aerated soils as Fe3+ ions, as Fe(II) will readily oxidise to Fe(III), which in turn forms Fe(III) hydroxides. However, the solubility of these hydroxides is highly dependent on pH. The reaction is:
Fe3+ + 3OH- ↔ Fe(OH)3
whereby iron(III) + hydroxide (OH-) ions are in an equilibrium with iron(III) hydroxide. At low pH the equilibrium shifts to the left, and more Fe(III) ions form. At high pH however, the formation (and precipitation) of solid iron(III) hydroxide is favoured, and the equilibrium shifts to the right. In fact, for every rise in pH of one unit, the activity of Fe(III) decreases one thousand-fold.
At pH 7 to 9, Fe(OH)2+, Fe(OH)3 and FeOH2+ form, with minimum solubility between pH 7.4 to 8.5. This means acid soils (pH less than 7) are higher in soluble iron than alkaline ones (pH above 7). Iron deficiencies for many plant species can arise in extremely alkaline soils. Acid-loving plants in particular cannot extract the iron they need in even mildly alkaline soils — these plants require larger amounts of iron than others, which are only available at a pH of 5.5 or even lower.
Waterlogged soils are low in oxygen, and in this environment Fe(III) reduces to Fe(II). This reduction is done by anaerobic bacteria, which use the electron obtained in respiration:
Fe(OH)3 + e- + 3H+ → Fe2+ + 3H2O
Fe(II) is usually found in higher concentrations than Fe(III) further down a soil profile where oxygen is less available, or even absent.
So far we’ve discussed inorganic Fe, that without carbon. However, iron — when soluble — readily forms organic complexes, or chelates, with organic matter particles. This has great importance in both the mobility of iron through soil, and the availability of iron to plants. Chelates in general make it possible for iron to leach through a soil profile, and specific chelates called siderophores are the most important bioavailable forms of iron for both plants and microbes. This is because inorganic iron in soil is often at levels far below the needs of plants, and the organic siderophores make more iron available.
Soil iron, whether in ionic or chelate form, must generally be reduced before it can be uptaken by roots. The rate of reduction is pH dependent and higher at low pH. Reducing the Fe(III) in chelate destabilises the complex and makes free Fe(II) available.
Some plant species are Fe efficient. When iron-stressed, these species are able to lower the pH of the soil immediately surrounding the roots (the rhizosphere) to increase the reducing capability of the roots for Fe(II) uptake. Fe inefficient species (mostly grasses) do not have this ability, but are generally less susceptible to lime induced chlorosis (more on this later). The Fe inefficient species can secrete their own versions of siderophores — phytosiderophores (plant siderophores) . These phytosiderophores mobilise the Fe in barely soluble iron compounds by forming very stable complexes with Fe(III) that are then uptaken by the roots. This process is independent of pH and occurs over a range from pH 4 to 8.
The two characteristics of Fe we’ve spent time on here — its ability to change oxidation state and its ability to form chelates — underpin its role in biochemistry. Iron has many roles in enzyme systems for which entire textbooks could be written, but just three include its role in the biosynthesis of chlorophyll enzymes, haemoglobin, and haem-containing enzymes. (Yes, plants contain haemoglobin!)
Fe competes for uptake with manganese (Mn2+), copper (Cu2+), calcium (Ca2+), magnesium (Mg2+), potassium (K+) and zinc (Zn2+) cations (positively-charged ions). These ions may induce an iron deficiency — Cu and Zn are particularly known to displace Fe from chelates to form their own.
Iron is not very mobile in plants, and deficiencies appear first in the younger plant parts, becoming chlorotic (yellowing due to lack of chlorophyll) whilst older parts remain green. The chlorosis is typically interveinal and in newly formed leaves displays a noticeable reticulated (forming a network) pattern of dark green veins against a lighter green or yellow background. The youngest leaves may even be completely white and devoid of chlorophyll.
While iron chlorosis as described above arises from an absolute deficiency of iron, lime induced chlorosis occurs specifically on calcareous soils (those containing calcium carbonate, or lime). Iron is present in these soils, but is not available to plants. These soils are of high carbonate (CO32-) concentrations, high calcium (Ca2+) concentrations, and high pH, and these conditions enable bicarbonate ions (HCO3-) to form:
CaCO3 + CO2 + H2O → Ca2+ + 2HCO3-
It is the bicarbonate ions that cause the deficiency. This ion has a physiological effect in that it both prevents uptake by and translocation within a plant.
The dissolution of calcium carbonate (CaCO3) to bicarbonate (HCO3-) requires carbon dioxide (CO2), which in the soil would come from root and microbial respiration. (Just like CO2 is a byproduct of our cellular respiration.) In a well-structured and aerated soil, CO2 would not accumulate at levels to cause HCO3- to accumulate. But in a poor soil that is very wet and doesn’t drain well, conditions are good for CO2 buildup and chlorosis. It’s best to not grow plants on these soils in the first place that aren’t lime-tolerant. It’s a losing battle to get the pH down to more manageable levels.
Iron toxicity is not common but too much iron can lead to the generation of hydroxyl radicals (•OH) which attack DNA, lipids, and protein. Symptoms include leaves with brown spots or leaves turning a ‘bronzed’ (really a deeper purple-brown) colour. High levels of iron can produce a manganese deficiency, which is often mistaken for an iron deficiency.