What are the environmental effects
of iron in water?
Iron is a dietary requirement for
most organisms, and plays an important role in natural processes in binary and
tertiary form. Oxidized tertiary iron cannot be applied by organisms freely,
except at very low pH values. Still, iron usually occurs in this generally
water insoluble form.
Adding soluble iron may rapidly increase productivity in oceanic surface layers. It might than play an important role in the carbon cycle. Iron is essential for nitrogen binding and nitrate reduction, and it may be a limiting factor for phytoplankton growth. Solubility in salt water is extremely low.
The iron cycle means reduction of tertiary iron by organic ligands (a process that is photo catalysed in surface waters), and oxidation of binary iron.
Iron forms chelation complexes that often play an important role in nature, such as haemoglobin, a red colouring agent in blood that binds and releases oxygen in breathing processes. Organisms take up higher amounts of binary iron than of tertiary iron, and uptake mainly depends on the degree of saturation of physical iron reserves.
Iron is often a limiting factor for water organisms in surface layers. When chelation ligands are absent, water insoluble tertiary iron hydroxides precipitate. This is not thought to be hazardous for aquatic life, because not much is known about hazards of water borne iron.
Mollusks have teeth of magnetite of goethite.
Green plants apply iron for energy transformation processes. Plants that are applied as animal feed may contain up to 1000 ppm of iron, but this amount is much lower in plants applied for human consumption. Generally plants contain between 20 and 300 ppm iron (dry mass), but lichens may consist up to 5.5% of iron. When soils contain little iron, or little water soluble iron, plants may experience growth problems. Plant uptake capacity strongly varies, and it does not only depend on soil iron concentrations, but also upon pH values, phosphate concentrations and competition between iron and other heavy metals. Limes soils are often iron deficit, even when sufficient amounts of iron are present. This is because of the generally high pH value, which leads to iron precipitation.
Iron usually occurs in soils in tertiary form, but in water saturated soils it is converted to binary iron, thereby enabling plant iron uptake. Plants may take up water insoluble iron compounds by releasing H+ ions, causing it to dissolve. Micro organisms release iron siderochrome, which can be directly taken up by plants.
Iron may be harmful to plants at feed concentrations of between 5 and 200 ppm. These cannot be found in nature under normal conditions, when low amounts of soil water are present.
A number of bacteria take up iron particles and convert them to magnetite, to apply this as a magnetic compass for orientation. Iron compounds may cause a much more serious environmental impact than the element itself. A number of LD50 values are known for rats (oral intake): iron (III) acetyl acetonate 1872 mg/kg, iron (II) chloride 984 mg/kg, and iron penta carbonyl 25 mg/kg.
There are four naturally occurring non-radioactive iron isotopes. There are eight instable iron isotopes.
Adding soluble iron may rapidly increase productivity in oceanic surface layers. It might than play an important role in the carbon cycle. Iron is essential for nitrogen binding and nitrate reduction, and it may be a limiting factor for phytoplankton growth. Solubility in salt water is extremely low.
The iron cycle means reduction of tertiary iron by organic ligands (a process that is photo catalysed in surface waters), and oxidation of binary iron.
Iron forms chelation complexes that often play an important role in nature, such as haemoglobin, a red colouring agent in blood that binds and releases oxygen in breathing processes. Organisms take up higher amounts of binary iron than of tertiary iron, and uptake mainly depends on the degree of saturation of physical iron reserves.
Iron is often a limiting factor for water organisms in surface layers. When chelation ligands are absent, water insoluble tertiary iron hydroxides precipitate. This is not thought to be hazardous for aquatic life, because not much is known about hazards of water borne iron.
Mollusks have teeth of magnetite of goethite.
Green plants apply iron for energy transformation processes. Plants that are applied as animal feed may contain up to 1000 ppm of iron, but this amount is much lower in plants applied for human consumption. Generally plants contain between 20 and 300 ppm iron (dry mass), but lichens may consist up to 5.5% of iron. When soils contain little iron, or little water soluble iron, plants may experience growth problems. Plant uptake capacity strongly varies, and it does not only depend on soil iron concentrations, but also upon pH values, phosphate concentrations and competition between iron and other heavy metals. Limes soils are often iron deficit, even when sufficient amounts of iron are present. This is because of the generally high pH value, which leads to iron precipitation.
Iron usually occurs in soils in tertiary form, but in water saturated soils it is converted to binary iron, thereby enabling plant iron uptake. Plants may take up water insoluble iron compounds by releasing H+ ions, causing it to dissolve. Micro organisms release iron siderochrome, which can be directly taken up by plants.
Iron may be harmful to plants at feed concentrations of between 5 and 200 ppm. These cannot be found in nature under normal conditions, when low amounts of soil water are present.
A number of bacteria take up iron particles and convert them to magnetite, to apply this as a magnetic compass for orientation. Iron compounds may cause a much more serious environmental impact than the element itself. A number of LD50 values are known for rats (oral intake): iron (III) acetyl acetonate 1872 mg/kg, iron (II) chloride 984 mg/kg, and iron penta carbonyl 25 mg/kg.
There are four naturally occurring non-radioactive iron isotopes. There are eight instable iron isotopes.
No comments:
Post a Comment