Is Fe2O3 acidic in water

The iron content in seawater is around 1-3 ppb, but varies greatly and is higher in the Atlantic than in the Pacific. River water generally contains 0.5-1 ppm and groundwater up to 100 ppm of this element.
Most algae have an iron concentration of 20-200 ppm, with some brown algae even accumulating up to 4000 ppm. The bioconcentration factor of algae compared to seawater is thus around 104-105. Marine fish contain around 10-90 ppm and oyster tissue around 195 ppm (values ​​are based on dry matter).
In dissolved form, iron is in the acidic and neutral pH range and under oxygen-rich conditions mainly as trivalent Fe (OH)2+(aq) present. Under reduced conditions it is more likely to occur in a bivalent form. In addition, iron is a component of many organic and inorganic chelates, which in most cases are readily soluble.

How and in which compounds does iron react with water?

While iron does not change significantly in pure water or in dry air, it rusts in the presence of water and oxygen, for example in moist air. The silvery color changes to brownish-red as the hydrogenated oxide is formed. Electrolytes dissolved in water accelerate the reaction, which looks like this:

4Fe + 3O2 + 6H2O -> 4Fe3+ + 12 (OH)- -> 4Fe (OH)3 or 4FeO (OH) + 4H2O

Most of the time, the rust layer that is created does not protect the iron from a further reaction, but dissolves so that more metal can oxidize. The electrolytes present are mostly iron (II) sulphate, which is released when attacked by atmospheric SO2 arises. In regions near the sea, salt particles in the air also play an important role.
Iron (III) hydroxide also often precipitates in natural waters.


Water solubility of iron and / or its compounds

Elemental iron is insoluble in water under normal circumstances. Numerous iron compounds also have this property. For example, naturally occurring iron oxide, just like iron hydroxide, iron carbide or iron pentacarbonyl, is insoluble in water. Some iron compounds have a higher solubility at a lower pH.
Other iron compounds, on the other hand, are more or less soluble in water. Iron carbonate has a water solubility of 60 mg / L, iron sulfide of 6 mg / L and iron vitriol of even 295 g / L. Many iron-containing chelate compounds are also readily soluble in water.
A distinction can usually be made between soluble Fe2+Compounds and generally insoluble Fe3+-Links. The latter are only soluble in very acidic solutions, but can, under certain circumstances, be reduced to Fe2+ and thus increase in solubility.

Solubility and how it can be influenced

How can iron get into water?

The most important minerals in which iron occurs naturally are magnetite, hematite, goethite, lepidocrocite and siderite. Through weathering processes, it also finds its way into bodies of water. Iron is contained in mineral water in the form of carbonate and is also found in drinking water. Fist-sized lumps of iron, manganese and small amounts of lime, silicon dioxide and organic compounds can often be found in deep sea areas.
Iron is of great commercial use and is produced annually in quantities of 500 million tons worldwide. In addition, around 300 million tons are recycled. This is probably also due to the fact that iron probably has more applications than any other metal. Above all, alloys ensure that the metal becomes more resistant to corrosion. In steel production, for example, carbon is added in various amounts. Iron alloys are finally processed into, for example, containers, cars, washing machines, bridges, buildings or even the smallest spring springs.
Iron compounds also serve as color pigments in glass and enamel production or are used in pharmaceutical products, the chemical industry, as iron fertilizer, in weed control, wood impregnation and photography.
Iron-containing "red mud", a waste product from aluminum extraction, was previously discharged into water. Today it is held back and is used as soil filling material.
Iron compounds are also used for precipitation reactions to remove water pollution in the so-called third purification stage.
The isotope 59Fe is used in medical diagnostics and nuclear physics research.

What environmental problems can water pollution with iron cause?

Iron is an essential element for almost all living beings and plays a major role in natural processes in its divalent and trivalent form. The oxidized trivalent form is usually difficult to use by organisms, except when the pH value is very low. Nevertheless, it occurs almost exclusively in this mostly insoluble form on earth.
Adding soluble iron could, for example, significantly increase productivity in the upper layers of the oceans. It could also play an important role in the carbon cycle. Since iron is essential for nitrogen fixation and nitrate reduction, it is likely a limiting factor in phytoplankton growth. The solubility in salt water is very low.
The iron cycle involves the reducing solution of trivalent iron by organic ligands, a process that can be photocatalyzed in surface water, and the oxidation of divalent iron by oxygen.
Iron forms chelates that often play an important role in nature, such as hemoglobin, the red blood pigment that can bind oxygen and release it again, which is of great importance for the respiratory processes. The uptake of bivalent iron in warm-blooded animals is significantly higher than that of the trivalent form and depends, among other things, on the degree of saturation of the iron stores in the body.
As far as aquatic organisms are concerned, iron often plays the role of a limiting factor in the upper layers of the seas. If there are no chelating agents in the water, iron (III) hydroxide, which is readily soluble at low pH values, precipitates and is generally unusable for organisms. However, it does not appear to be very harmful to aquatic life either, as little is generally known about the harmfulness of iron to aquatic organisms.
Mollusks make their teeth out of magnetite or goethite.
Green plants use iron, among other things, for processes in which they have to transform energy. Forage plants can even have an iron content of up to 1000 ppm; this value is much lower for plant-based foods for humans. In general, the concentration of iron in plants is around 20-300 ppm (based on dry matter), whereby lichens can even consist of 5.5% of this element. Growth problems can be observed when the soil does not contain enough iron or it is only available in a highly insoluble form.
The absorption capacity of plants varies greatly and is not only dependent on the availability of iron in the soil, but also on the pH value, the phosphate content and the competition with other heavy metals in the soil. Iron deficiency symptoms often occur on very calcareous soils, even if there is enough iron. This is because the pH of the soil increases due to the lime content and there is no longer enough iron in dissolved form.
Usually iron occurs as iron (III) in the soil, but is converted into iron (II) in water-saturated soils and can thus be better absorbed by the plant roots. Plants can absorb the element from poorly soluble compounds, for example by using H+-Release ions, which brings it into solution. Microorganisms also release iron-containing siderochrome as a metabolic product, which higher plants can absorb directly.
Regarding the toxicity of iron to plants, harmful effects have been observed at concentrations of 5-200 ppm in nutrient solutions. However, these rarely occur in nature if there is no impact from backwater.
Some bacteria have a special use for iron. They pick up iron particles in order to convert them into magnetite and in turn to be able to use this material as a magnetic compass for their sense of direction.
Iron compounds can have a significantly greater negative impact on the environment than the element itself. So are some LDs50-Values ​​known that indicate the concentration of a substance at which half of a population dies. For example, the LD50-Value of iron (III) acetylacetonate when taken orally by the rat 1872 mg / kg, that of iron (II) chloride 984 mg / kg and that of iron pentacarbonyl even only 25 mg / kg.
There are naturally four iron isotopes that are not radioactive. However, there are now eight unstable isotopes as well.

What health effects can iron cause in water?

The total amount of iron in the human body is around 4 g, of which around 70% is in the red blood pigment. As for almost all organisms, the element is also essential for humans. Men need around 7 mg of iron per day, while women, because of menstruation, need as much as 11 mg. With a normal diet, this amount is usually not a problem, of which approximately 25% is ultimately absorbed, depending on the supplies in the body. This rate can be increased with vitamin C, as it reduces trivalent to divalent iron, while it is rather lower with phosphates and phytates.
In food, iron is either in a divalent form, the ferroform, bound to hemoglobin and myoglobin as heme iron, or in trivalent ferric form as non-heme iron. Especially the mostly animal ferroform can be easily absorbed by the body.
As the central atom of the red blood pigment hemoglobin, iron binds oxygen to itself and transports it from the lungs to all other parts of the body. From there it in turn transports CO2 to the lungs where it can be exhaled. Iron is also required for oxygen storage. It is also a component of various enzymes and is therefore involved in DNA synthesis, for example. Ordinary brain functions are also dependent on iron. In addition, iron in the body is strongly bound to transferrin, which is responsible for exchanging the metal between cells. This substance has a strong antibiotic effect because it withholds the vital substance from unwanted bacteria. As the bacterial infestation increases, the body therefore also produces significantly more transferrin.
The excess iron stores are stored in the liver, but bone marrow is also very rich in iron, as this is where hemoglobin is formed.
Iron deficiency results in anemia (anemia), which manifests itself in fatigue, headaches and decreased performance. The immune system is also weakened. In children it has a negative influence on the mental development of these, causing excessive irritability and learning difficulties.
In the case of iron deficiency, which occurs especially in small children, pregnant women and menstruating women, soluble iron (II) salts are often given.
If the iron intake is too high, which occurs, for example, in hemochromatosis patients, iron is stored in the pancreas, liver, spleen and heart, which can lead to organ damage. In healthy people, iron, which is hardly toxic, actually has no negative consequences and overloading occurs only very rarely in drinking water with an iron content of over 200 ppm.
Iron compounds can have significantly more serious health effects than the relatively harmless iron itself. Soluble Fe2 + compounds, such as FeCl, in particular have2 or FeSO4, have a toxic effect from an ingestion of about 200 mg and are fatal to adult humans in a dose of 10-50 g. Some iron-containing chelates can also be very dangerous, and iron carbonyls, such as the nerve agent iron pentacarbonyl, are known to be toxic. Iron dust can cause lung disease.

What water purification technologies can be used to remove iron?

Iron removal is particularly common in drinking water treatment, as well water, for example, can contain many iron ions that affect the color, cloudiness and taste of the water. However, iron can also be found in virtually all wastewater.
In order to remove iron from water, the divalent iron can be oxidized to the trivalent form, whereupon iron hydroxide flakes are formed through hydrolysis, which can be filtered off with the help of sand filters, for example. The oxidation can be brought about by oxygen or the use of oxidizing agents such as chlorine or potassium permanganate. The speed of this reaction depends on the pH value and takes place more slowly in an acidic environment than in a basic one. To speed up the reaction at low pH levels, the water can be aerated to remove carbon dioxide and raise the pH of the water. The overall reaction is acid-forming and therefore slows itself down. Often iron is reduced along with manganese.
It is also possible to use ion exchangers to remove iron, but they are more suitable for removing iron traces in drinking and process water than for large concentrations of iron.
Iron compounds themselves are also used in water treatment, mostly as precipitants, for example in the form of iron sulfate for phosphate precipitation.

The drinking water standards of the EU and Germany specify a maximum iron concentration of 0.2 mg / L.

Comparison of drinking water standards

References


To the periodic table of the elements

To the overview of the elements and water