Iron plays an important role in biology, forming complexes with molecular oxygen in hemoglobin and myoglobin; these two compounds are common oxygen transport proteins in vertebrates. Iron is also the metal at the active site of many important redox enzymes dealing with cellular respiration and oxidation and reduction in plants and animals. A human male of average height has about 4 grams of iron in his body, a female about 3.5 grams. This iron is distributed throughout the body in hemoglobin, tissues, muscles, bone marrow, blood proteins, enzymes, ferritin, hemosiderin, and transport in plasma.
Biological and pathological role
Main article: Human iron metabolism
Iron is involved in numerous biological processes. It is the most important transition metal in all living organisms. Iron-proteins are found in all living organisms: archaeans, bacteria and eukaryotes, including humans. For example, the color of blood is due to hemoglobin, an iron-containing protein. As illustrated by hemoglobin, iron is often bound to cofactors, such as hemes, which are non-protein compounds, often involving metal ions, that are required for a protein's biological activity to happen. The iron-sulfur clusters are pervasive and include nitrogenase, the enzymes responsible for biological nitrogen fixation. The main roles of iron-containing proteins are the transport and storage of oxygen, as well as the transfer of electrons.
Iron is a necessary trace element found in nearly all living organisms. Iron-containing enzymes and proteins, often containing heme prosthetic groups, participate in many biological oxidations and in transport. Examples of proteins found in higher organisms include hemoglobin, cytochrome (see high-valent iron), and catalase. The average adult human contains about 0.005% body weight of iron, or about four grams, of which three quarters is in hemoglobin – a level that remains constant despite only about one milligram of iron being absorbed each day, because the human body recycles its hemoglobin for the iron content.
Iron acquisition poses a problem for aerobic organisms because ferric iron is poorly soluble near neutral pH. Thus, these organisms have developed means to absorb iron as complexes, sometimes taking up ferrous iron before oxidising it back to ferric iron. In particular, bacteria have evolved very high-affinity sequestering agents called siderophores.
After uptake in human cells, iron storage is carefully regulated; iron ions are never "free". This is because free iron ions have a high potential for biological toxicity. A major component of this regulation is the protein transferrin, which binds iron ions absorbed from the duodenum and carries it in the blood to cells. Transferrin contains Fe3+ in the middle of a distorted octahedron, bonded to one nitrogen, three oxygens and a chelating carbonate anion that traps the Fe3+ ion: it has such a high stability constant that it is very effective at taking up Fe3+ ions even from the most stable complexes. At the bone marrow, transferrin is reduced from Fe3+ and Fe2+and stored as ferritin to be incorporated into hemoglobin.
The most commonly known and studied bioinorganic iron compounds (biological iron molecules) are the heme proteins: examples are hemoglobin, myoglobin, and cytochrome P450. These compounds participate in transporting gases, building enzymes, and transferring electrons. Metalloproteins are a group of proteins with metal ion cofactors. Some examples of iron metalloproteins are ferritin and rubredoxin. Many enzymes vital to life contain iron, such as catalase, lipoxygenases, and IRE-BP.
Hemoglobin is an oxygen carrier that occurs in red blood cells and contributes their color, transporting oxygen in the arteries from the lungs to the muscles where it is transferred to myoglobin, which stores it until it is needed for the metabolic oxidation of glucose, which generates energy. Here the hemoglobin binds to carbon dioxide, produced when glucose is oxidized, which is transported through the veins by hemoglobin (predominantly as bicarbonate anions) back to the lungs where it is exhaled. In hemoglobin, the iron is in one of four heme groups and has six possible coordination sites; four are occupied by nitrogen atoms in a porphyrin ring, the fifth by an imidazole nitrogen in a histidine residue of one of the protein chains attached to the heme group, and the sixth is reserved for the oxygen molecule it can reversibly bind to. When hemoglobin is not attached to oxygen (and is then called deoxyhemoglobin), the Fe2+ ion at the center of the heme group (in the hydrophobic protein interior) is in a high-spin configuration. It is thus too large to fit inside the porphyrin ring, which bends instead into a dome with the Fe2+ ion about 55 picometers above it. In this configuration, the sixth coordination site reserved for the oxygen is blocked by another histidine residue. When deoxyhemoglobin picks up an oxygen molecule, this histidine residue moves away and returns once the oxygen is securely attached to form a hydrogen bond with it. This results in the Fe2+ ion switching to a low-spin configuration, resulting in a 20% decrease in ionic radius so that now it can fit into the porphyrin ring, which becomes planar. (Additionally, this hydrogen bonding results in the tilting of the oxygen molecule, resulting in a Fe–O–O bond angle of around 120° that avoids the formation of Fe–O–Fe or Fe–O2–Fe bridges that would lead to electron transfer, the oxidation of Fe2+ to Fe3+, and the destruction of hemoglobin.) This results in a movement of all the protein chains that leads to the other subunits of hemoglobin changing shape to a form with larger oxygen affinity. Thus, when deoxyhemoglobin takes up oxygen, its affinity for more oxygen increases, and vice versa. Myoglobin, on the other hand, contains only one heme group and hence this cooperative effect cannot occur. Thus, while hemoglobin is almost saturated with oxygen in the high partial pressures of oxygen found in the lungs, its affinity for oxygen is much lower than myoglobin in the low partial pressures of oxygen found in muscle tissue, resulting in oxygen transfer. This is further enhanced by the concomitant Bohr effect (named after Christian Bohr, the father of Niels Bohr), in which lowered pH (as occurs when carbon dioxide is released in the muscles) further lowers the oxygen affinity of hemoglobin.
Carbon monoxide and phosphorus trifluoride are poisonous to humans because they bind to hemoglobin similarly to oxygen, but with much more strength, so that oxygen can no longer be transported throughout the body. This effect also plays a minor role in the toxicity of cyanide, but there the major effect is by far its interference with the proper functioning of the electron transport protein cytochrome a. The cytochrome proteins also involve heme groups and are involved in the metabolic oxidation of glucose by oxygen. The sixth coordination site is then occupied by either another imidazole nitrogen or a methionine sulfur, so that these proteins are largely inert to oxygen – with the exception of cytochrome a, which bonds directly to oxygen and thus is very easily poisoned by cyanide. Here, the electron transfer takes place as the iron remains in low spin but changes between the +2 and +3 oxidation states. Since the reduction potential of each step is slightly greater than the previous one, the energy is released step-by-step and can thus be stored in adenosine triphosphate. Cytochrome a is slightly distinct, as it occurs at the mitochondrial membrane, binds directly to oxygen, and transports protons as well as electrons, as follows:
- 4 Cytc2+ + O2 + 8H+
inside → 4 Cytc3+ + 2 H2O + 4H+
Although the heme proteins are the most important class of iron-containing proteins, the iron-sulfur proteins are also very important, being involved in electron transfer, which is possible since iron can exist stably in either the +2 or +3 oxidation states. These have one, two, four, or eight iron atoms that are each approximately tetrahedrally coordinated to four sulfur atoms; because of this tetrahedral coordination, they always have high-spin iron. The simplest of such compounds is rubredoxin, which has only one iron atom coordinated to four sulfur atoms from cysteine residues in the surrounding peptide chains. Another important class of iron-sulfur proteins is the ferredoxins, which have multiple iron atoms. Transferrin does not belong to either of these classes.
Health and diet
Iron is pervasive, but particularly rich sources of dietary iron include red meat, lentils, beans, poultry, fish, leaf vegetables, watercress, tofu, chickpeas, black-eyed peas, and blackstrap molasses. Bread and breakfast cereals are sometimes specifically fortified with iron. Iron in low amounts is found in molasses, teff, and farina.
Iron provided by dietary supplements is often found as iron(II) fumarate, although iron(II) sulfate is cheaper and is absorbed equally well. Elemental iron, or reduced iron, despite being absorbed at only one-third to two-thirds the efficiency (relative to iron sulfate), is often added to foods such as breakfast cereals or enriched wheat flour. Iron is most available to the body when chelated to amino acids and is also available for use as a common iron supplement. Glycine, the cheapest and most common amino acid is most often used to produce iron glycinate supplements. The Recommended Dietary Allowance (RDA) for iron varies considerably depending on age, sex, and source of dietary iron: for example, heme-based iron has higher bioavailability.
Dietary reference intake
The Food and Nutrition Board of the U.S. Institute of Medicine updated Estimated Average Requirements (EARs) and Recommended Dietary Allowances (RDAs) for iron in 2001. The current EAR for iron for women ages 14-18 is 7.9 mg/day, 8.1 for ages 19-50 and 5.0 thereafter (post menopause). For men the EAR is 6.0 mg/day for ages 19 and up. The RDA is 15.0 mg/day for women ages 15-18, 18.0 for 19-50 and 8.0 thereafter. For men, 8.0 mg/day for ages 19 and up. RDAs are higher than EARs so as to identify amounts that will cover people with higher than average requirements. RDA for pregnancy equals 27 mg/day. RDA for lactation equals 9 mg/day. For children ages 1–3 years 7 mg/day, 10 for ages 4-8 and 8 for ages 9-13. As for safety, the Food and Nutrition Board also sets Tolerable Upper Intake Levels (known as ULs) for vitamins and minerals when evidence is sufficient. In the case of iron the UL is set at 45 mg/day. Collectively the EARs, RDAs and ULs are referred to as Dietary Reference Intakes. The European Food Safety Authority reviewed the same safety question did not establish a UL.
For U.S. food and dietary supplement labeling purposes the amount in a serving is expressed as a percent of Daily Value (%DV). For iron labeling purposes 100% of the Daily Value was 18.0 mg, and as of May 2016 remained unchanged at 18.0 mg. Food and supplement companies have until July 28, 2018 to comply with the change. A table of the pre-change adult Daily Values is provided at Reference Daily Intake.
Main article: Iron overload
Iron uptake is tightly regulated by the human body, which has no regulated physiological means of excreting iron. Only small amounts of iron are lost daily due to mucosal and skin epithelial cell sloughing, so control of iron levels is primarily accomplished by regulating uptake. Regulation of iron uptake is impaired in some people as a result of a genetic defect that maps to the HLA-H gene region on chromosome 6 and leads to abnormally low levels of hepcidin, a key regulator of the entry of iron into the circulatory system in mammals. In these people, excessive iron intake can result in iron overload disorders, known medically as hemochromatosis. Many people have an undiagnosed genetic susceptibility to iron overload, and are not aware of a family history of the problem. For this reason, people should not take iron supplements unless they suffer from iron deficiency and have consulted a doctor. Hemochromatosis is estimated to be the a cause of 0.3 to 0.8% of all metabolic diseases of Caucasians.
Overdoses of ingested iron can cause excessive levels of free iron in the blood. High blood levels of free ferrous iron react with peroxides to produce highly reactive free radicals that can damage DNA, proteins, lipids, and other cellular components. Iron toxicity occurs when the cell contains free iron, which generally occurs when iron levels exceed the availability of transferrin to bind the iron. Damage to the cells of the gastrointestinal tract can also prevent them from regulating iron absorption, leading to further increases in blood levels. Iron typically damages cells in the heart, liver and elsewhere, causing adverse effects that include coma, metabolic acidosis, shock, liver failure, coagulopathy, adult respiratory distress syndrome, long-term organ damage, and even death. Humans experience iron toxicity when the iron exceeds 20 milligrams for every kilogram of body mass; 60 milligrams per kilogram is considered a lethal dose. Overconsumption of iron, often the result of children eating large quantities of ferrous sulfate tablets intended for adult consumption, is one of the most common toxicological causes of death in children under six. The Dietary Reference Intake (DRI) sets the Tolerable Upper Intake Level (UL) for adults at 45 mg/day. For children under fourteen years old the UL is 40 mg/day.
Main article: Iron deficiency
Iron deficiency is the most common nutritional deficiency in the world. When loss of iron is not adequately compensated by adequate dietary iron intake, a state of latent iron deficiency occurs, which over time leads to iron-deficiency anemia if left untreated, which is characterised by an insufficient number of red blood cells and an insufficient amount of hemoglobin. Children, pre-menopausal women (women of child-bearing age), and people with poor diet are most susceptible to the disease. Most cases of iron-deficiency anemia are mild, but if not treated can cause problems like fast or irregular heartbeat, complications during pregnancy, and delayed growth in infants and children.