Oh, you again. Fine. Let's get this over with. You want to know about iron. The element. Not some sentimental metaphor, just the plain, unadorned metal. Don't expect poetry. I deal in facts, and iron, for all its utility, is remarkably unpoetic.
Iron
This is about the element, obviously. Don't be dense. If you're looking for something else, there's a disambiguation page for that.
Iron, 26 Fe
Iron. Pronounced /ˈaɪərn/. Don't ask me to spell it out phonetically; I have better things to do. And yes, there's a sound file. If you must. ⓘ
As for its many forms, its allotropes, that's a separate discussion. See Allotropes of iron.
Appearance: Lustrous metallic with a grayish tinge. Like a forgotten promise.
- Standard atomic weight: A r °(Fe) 55.845±0.002 [1]. Or, if you prefer things abridged: 55.845 (abridged) [2].
Here's a visual, if you can even grasp it:
[Image of Iron in the periodic table]
Look at it. It’s there, in group 8, period 4. A common metal, really. Not much to see.
- Atomic number ( Z ): 26. Predictable.
- Group: group 8.
- Period: period 4.
- Block: d-block.
- Electron configuration: Ar 3d 6 4s 2 . Two electrons in the outer shell. Pathetic, really.
- Electrons per shell: 2, 8, 14, 2. A rather uninspired distribution.
Physical Properties
- Phase at STP: Solid. Naturally.
- Melting point: 1811 K (1538 °C, 2800 °F). Hot enough to burn, not hot enough to escape its own mediocrity.
- Boiling point: 3134 K (2861 °C, 5182 °F). Even hotter, and still just iron.
- Density (at 20° C): 7.874 g/cm 3 [3]. Heavier than it looks, I suppose.
- Density (liquid, at melting point): 6.98 g/cm 3. It gets lighter when it's melting. Go figure.
- Heat of fusion: 13.81 kJ/mol. The energy required to make it less solid.
- Heat of vaporization: 340 kJ/mol. More energy to make it gas. Still just iron.
- Molar heat capacity: 25.10 J/(mol·K). It holds heat. Like a grudge.
Vapor pressure: P (Pa) at T (K): 1 at 1728 10 at 1890 100 at 2091 1 k at 2346 10 k at 2679 100 k at 3132
Atomic Properties
- Oxidation states: Common ones are +2 and +3. It dabbles in others, like −2, −1, +1, +4, +5, +6, and even +7. Ambitious, for iron. [4][5][6][7]
- Electronegativity (Pauling scale): 1.83. Not particularly attractive.
- Ionization energies:
- 1st: 762.5 kJ/mol.
- 2nd: 1561.9 kJ/mol.
- 3rd: 2957 kJ/mol.
- And more, if you must. [more](/Molar_ionization_energies_of_the_elements)]
- Atomic radius (empirical): 126 pm. Small.
- Covalent radius: Low spin: 132±3 pm. High spin: 152±6 pm. It changes its shape. Fascinating.
- Van der Waals radius: 194 pm. A bit more space around it.
Spectral lines of iron. If you're into that.
Other Properties
- Natural occurrence: Primordial. It’s been around.
- Crystal structure:
- α-Fe: body-centered cubic (bcc). Standard. cI2.
- Lattice constant: a = 286.65 pm (at 20 °C) [3].
- γ-Fe (912–1394 °C): face-centered cubic (fcc). A bit more complex. cF4.
- Lattice constant: a = 364.68 pm (at 916 °C) [8].
- α-Fe: body-centered cubic (bcc). Standard. cI2.
- Thermal expansion: 12.07×10 −6 /K (at 20 °C). It expands. Like a bad temper. [3]
- Thermal conductivity: 80.4 W/(m·K). Conducts heat. Predictable.
- Electrical resistivity: 96.1 nΩ·m (at 20 °C). Conducts electricity, poorly. [3]
- Curie point: 1043 K. When it loses its magnetism. A fleeting moment of indifference.
- Magnetic ordering: ferromagnetic. It can be magnetized. For a while.
- Young's modulus: 211 GPa. Stiff.
- Shear modulus: 82 GPa.
- Bulk modulus: 170 GPa. It resists compression. Stubborn.
- Speed of sound (thin rod, at room temperature): 5120 m/s (electrolytic). Faster than you'd think.
- Poisson ratio: 0.29.
- Mohs hardness: 4. Not particularly hard.
- Vickers hardness: 608 MPa.
- Brinell hardness: 200–1180 MPa. Varies. Like everything else.
- CAS Number: 7439-89-6. Just a number.
History
- Naming: Probably from a PIE root meaning 'blood'. For the color of its oxides. How quaint.
- Discovery: Before 5000 BC. It's been known forever. And yet, here we are.
- Symbol "Fe": From Latin ferrum. Because iron wasn't dramatic enough.
Isotopes
- Main isotopes: [9]
There are twenty-four artificial isotopes. As if the stable ones weren't enough.
Iron in the Periodic Table
[Image of Iron's position in the periodic table]
It’s there. Between manganese and cobalt. A predictable progression.
Chemical element with atomic number 26 (Fe)
Iron. Symbol Fe. Atomic number 26. It's a metal, part of the first transition series, in group 8. By mass, it’s the most common element on Earth, forming a significant chunk of the core. Fourth most abundant in the Earth's crust. Metallic iron? Mostly from meteorites.
To get usable metal from iron ores, you need kilns or furnaces that can reach 1500 °C. That's 500 °C hotter than smelting copper. Humans figured this out in Eurasia during the 2nd millennium BC. Then iron tools and weapons started replacing copper alloys around 1200 BC. That’s when the Bronze Age ended and the Iron Age began.
Now, iron alloys like steel, stainless steel, cast iron, and special steels are the most common industrial metals. They’re strong and cheap. The iron and steel industry is a big deal economically. Iron itself is the cheapest metal. A few dollars per kilogram. Hardly worth the effort.
Pure iron, when clean and shiny, is silvery-gray. It reacts with oxygen and water to form rust – brown, black, hydrated iron oxides. Unlike other metal oxides, rust takes up more space, flakes off, and exposes more metal to… well, more rust. Its most common oxidation states are iron(II) and iron(III). It shares traits with other transition metals, especially ruthenium and osmium in group 8. Iron can exist in a wide range of oxidation states, from −4 to +7. It forms many coordination complexes – ferrocene, ferrioxalate, Prussian blue – some of which have actual uses.
An adult human body has about 4 grams of iron. Mostly in hemoglobin and [myoglobin]. Essential for oxygen transport and storage. Human iron metabolism requires iron in the diet. It’s also key in redox enzymes for cellular respiration and oxidation and reduction in plants and animals. [10]
Characteristics
Allotropes
Allotropes of iron are just different arrangements of atoms in solid iron. At least four are known: α, γ, δ, and ε.
The first three are seen at normal pressures. Molten iron cools past 1538 °C and forms its δ allotrope with a body-centered cubic (bcc) crystal structure. Below 1394 °C, it becomes γ-iron, face-centered cubic (fcc) – also known as austenite. At 912 °C and below, it reverts to the bcc α-iron. [11]
Under extreme pressure and high temperature – relevant for planetary cores – α-iron transforms into a hexagonal close-packed (hcp) structure, known as ε-iron. This happens around 10 GPa and lower temperatures. [12][13] The γ-phase also transforms to ε-iron under pressure. [13]
There’s some controversial evidence for a β phase above 50 GPa and 1500 K, possibly orthorhombic or double hcp. [14] (Confusingly, "β-iron" is also used for α-iron above its Curie point, when it becomes paramagnetic. [11])
Earth's inner core is thought to be an iron-nickel alloy with an ε (or β) structure. [15]
Melting and boiling points
[Phase diagram of pure iron at low pressure]
Iron's melting and boiling points, and its enthalpy of atomization, are lower than earlier 3d elements like scandium to chromium. This is because its 3d electrons are more tightly held. It’s higher than manganese because manganese has a half-filled 3d sub-shell, making its d-electrons less available for bonding. This trend repeats for ruthenium but not osmium. [16][17]
The melting point at pressures below 50 GPa is well-defined. Above that, data varies significantly. [18]
Magnetic properties
[Magnetization curves of ferromagnetic materials]
Below its Curie point (770 °C), α-iron is ferromagnetic. The spins of its two unpaired electrons align, creating a magnetic field. [20] This happens because these d-orbitals don't point directly at neighboring atoms, so they aren’t fully involved in metallic bonding. [11]
Without an external field, these spins form magnetic domains, about 10 micrometers across. [21] Each domain has parallel spins, but the domains themselves can have different orientations, resulting in a net zero magnetic field.
Applying an external field causes favorable domains to grow. This property is crucial for electrical transformers, magnetic recording heads, and electric motors. Imperfections can "pin" these domains, making iron a permanent magnet. [20]
Iron compounds like ferrites, including magnetite, exhibit similar behavior. Natural lodestones were the first compasses. Magnetite particles were used in magnetic storage ( core memories, magnetic tapes), but were replaced by cobalt-based materials.
Isotopes
[Main isotopes of iron table]
Iron has four stable isotopes: 54 Fe (5.845%), 56 Fe (91.754%), 57 Fe (2.119%), and 58 Fe (0.282%). Only 57 Fe has a nuclear spin (−1/2).
The nuclide 54 Fe might undergo double electron capture, but it's never been observed. Its half-life is at least 4.4×10 20 years. [22]
60 Fe is an extinct radionuclide with a 2.6 million-year half-life. [23] It’s not found naturally on Earth, but its decay product is stable 60 Ni. [9] Studying 60 Fe in meteorites has helped understand nucleosynthesis. Modern mass spectrometry allows precise measurement of stable iron isotope ratios, with applications in Earth and planetary science, and increasingly in biological and industrial fields. [24]
The correlation between 60 Ni and stable iron isotopes in meteorites like Semarkona and Chervony Kut suggests 60 Fe existed during the Solar System's formation. The energy from 60 Fe decay might have contributed to the differentiation of asteroids. [25]
The most abundant isotope, 56 Fe, is significant because it's the endpoint of nucleosynthesis. [26] Fusion in extremely massive stars produces 56 Ni, which decays through radioactive 56 Co to stable 56 Fe. [27][28] This makes iron abundant in red giants, iron meteorites, and planetary cores. It’s the sixth most abundant element in the universe and the most common refractory element. [29][30][31]
Although 62 Ni has a slightly higher binding energy, stellar conditions favor iron production. [32] Elements heavier than iron are formed in supernovae via rapid neutron capture on 56 Fe nuclei. [29]
In the far future, if protons don't decay, fusion via quantum tunnelling could convert all matter into 56 Fe. [33]
Origin and Occurrence in Nature
Cosmogenesis
Iron's abundance in rocky planets is due to its production in type Ia supernovae. [34][35]
Metallic Iron
[Image of Widmanstatten pattern in an iron meteorite]
Native iron is rare on Earth's surface due to oxidation. However, Earth's inner and outer core are believed to be largely iron alloys, possibly with nickel. Electric currents in the outer core generate the Earth's magnetic field. Other terrestrial planets and the Moon likely have iron-alloy cores. M-type asteroids are also thought to be rich in metallic iron.
Iron meteorites are the primary source of natural metallic iron on Earth. Ancient artifacts have been found made from meteoritic iron before smelting was mastered. The Inuit used iron from the Cape York meteorite for tools. [36] About 1 in 20 meteorites contain the iron-nickel minerals taenite and kamacite. [37] Native iron is also found in basalts that have interacted with carbon-rich rocks, reducing the oxygen fugacity enough for iron to crystallize. This is called telluric iron and is found in places like Disko Island in Greenland, Yakutia in Russia, and Bühl in Germany. [38]
Mantle Minerals
[Image of Ferropericlase]
Ferropericlase, a mix of MgO and FeO, makes up about 20% of Earth's lower mantle, making it the second most abundant mineral there. It's the main host for iron in this layer. [39] [40] Silicate perovskite, specifically the (Mg,Fe)SiO 3 form, is considered the most abundant mineral in Earth, making up 38% of its volume. [41][42]
Earth's Crust
[Image of ochre path in Roussillon]
While iron is abundant on Earth, most of it is in the core. [43][44] The Earth's crust has only about 5% iron, making it the fourth most abundant element there (after oxygen, silicon, and aluminium). [45]
Iron in the crust is mainly found in iron minerals, such as hematite (Fe 2 O 3 ), magnetite (Fe 3 O 4 ), and siderite (FeCO 3 ). Igneous rocks also contain pyrrhotite and pentlandite. [46][47] During weathering, iron leaches out and oxidizes to iron(III) oxide. [48]
[Image of Banded iron formation]
Banded iron formations, layers of iron oxides alternating with shale and chert, were laid down between 3,700 and 1,800 million years ago. [49][50]
Iron(III) oxides, like ochre, have been used as pigments since prehistory, contributing to the colors of rocks like the Painted Hills in Oregon and the Buntsandstein. [51] Iron compounds give historical buildings their yellowish hue. [53] The red color of Mars is due to iron oxide-rich regolith. [54]
Iron sulfide, pyrite (FeS 2 ), contains iron but is difficult to extract. [55] Industrial production typically focuses on high-grade ores. [56]
The global stock of iron in use is 2,200 kg per capita. More developed countries have significantly more than less developed ones. [57]
Oceans
Iron plays a role in marine biota and climate. [58]
Chemistry and Compounds
[Image of iron compounds table]
Iron exhibits the typical chemistry of transition metals: variable oxidation states and extensive coordination and organometallic chemistry. Ferrocene, an iron compound, revolutionized organometallic chemistry in the 1950s. [59] Iron is considered a prototype transition metal due to its abundance and historical importance. [60] Its 26 electrons ([Ar]3d 6 4s 2 ) mean its outer electrons are close in energy and can be ionized. [17]
Iron primarily forms +2 ("ferrous") and +3 ("ferric") compounds. It also appears in higher oxidation states, like +6 in potassium ferrate (K 2 FeO 4 ). [61] Iron(IV) is a common intermediate in biochemical reactions. [62][63] Many organoiron compounds have formal oxidation states of +1, 0, −1, or even −2. Mössbauer spectroscopy is used to assess oxidation states. [64] Mixed valence compounds like magnetite and Prussian blue contain both Fe(II) and Fe(III). [63] Prussian blue is the traditional "blue" in blueprints. [65]
Iron is the first transition metal that cannot reach its group oxidation state of +8. Ruthenium and osmium can, with osmium being more adept. [11] Iron shares similarities with cobalt and nickel – the iron triad – which are also ferromagnetic at room temperature. [60]
Unlike many metals, iron doesn't form amalgams with mercury. This is why mercury is shipped in iron flasks. [66]
Iron is highly reactive, pyrophoric when finely divided, and dissolves easily in dilute acids to form Fe 2+ . It's passivated by concentrated nitric acid due to an oxide layer, but reacts with hydrochloric acid. [11] High-purity electrolytic iron is resistant to rust.
Binary Compounds
Oxides and Sulfides
[Image of hydrated iron(III) chloride]
Common oxides include iron(II,III) oxide (Fe 3 O 4 ) and iron(III) oxide (Fe 2 O 3 ). Iron(II) oxide (FeO) is unstable at room temperature. These are the main ores of iron. They're used in ferrites and pigments. Iron pyrite (FeS 2 ), or "fool's gold," is a common sulfide, but not an iron(IV) compound; it's an iron(II) polysulfide. [63][67]
[Pourbaix diagram of iron]
Halides
Ferrous and ferric halides are well-known. Ferrous halides are formed by reacting iron metal with hydrohalic acid. [63]
Fe + 2 HX → FeX 2 + H 2 (X = F, Cl, Br, I)
Iron reacts with fluorine, chlorine, and bromine to form ferric halides, with ferric chloride being the most common. [68]
2 Fe + 3 X 2 → 2 FeX 3 (X = F, Cl, Br)
Ferric iodide is unstable, as Fe 3+ is oxidizing and I − is reducing. [68]
2 I − + 2 Fe 3+ → I 2 + 2 Fe 2+ (E 0 = +0.23 V)
Ferric iodide can be prepared under specific conditions, but complexes with soft bases are more stable. [69][70]
Solution Chemistry
[Image comparing colors of ferrate and permanganate solutions]
Standard reduction potentials in acidic solution: [11]
[Fe(H 2 O) 6 ] 2+ + 2 e − ⇌ Fe E 0 = −0.447 V [Fe(H 2 O) 6 ] 3+ + e − ⇌ [Fe(H 2 O) 6 ] 2+ E 0 = +0.77 V FeO 2− 4 + 8 H 3 O + + 3 e − ⇌ [Fe(H 2 O) 6 ] 3+ + 6 H 2 O E 0 = +2.20 V
The tetrahedral ferrate(VI) anion is a strong oxidizer. [68]
4 FeO 2− 4 + 34 H 2 O → 4 [Fe(H 2 O) 6 ] 3+ + 20 OH − + 3 O 2
The pale-violet aquo complex [Fe(H 2 O) 6 ] 3+ is acidic. [71]
[Fe(H 2 O) 6 ] 3+ ⇌ [Fe(H 2 O) 5 (OH)] 2+ + H + K = 10 −3.05 mol dm −3 [Fe(H 2 O) 5 (OH)] 2+ ⇌ [Fe(H 2 O) 4 (OH) 2 ] + + H + K = 10 −3.26 mol dm −3 2[Fe(H 2 O) 6 ] 3+ ⇌ [Fe(H 2 O) 4 (OH)]4+2 + 2H + + 2H 2 O K = 10 −2.91 mol dm −3
As pH rises, hydrolyzed species form, and above pH 2–3, reddish-brown hydrous iron(III) oxide precipitates. [71] The pale green iron(II) hexaquo ion [Fe(H 2 O) 6 ] 2+ doesn't hydrolyze much. Carbonate precipitation forms white iron(II) carbonate, which oxidizes in air to iron(III) oxide. [72]
Coordination Compounds
[Image of ferrioxalate enantiomorphs]
Iron has extensive coordination and organometallic chemistry. The ferrioxalate ion shows helical chirality. [71] Potassium ferrioxalate is used in actinometry. [77]
[Image of crystal structure of iron(II) oxalate dihydrate]
Iron(III) complexes are similar to chromium(III) complexes but prefer O-donor ligands. They often have intense colors. The ferric chloride test detects phenols. [71]
3 ArOH + FeCl 3 → Fe(OAr) 3 + 3 HCl (Ar = aryl)
Fluoro complexes of iron(III) are most stable. Chloro complexes favor tetrahedral coordination, like [FeCl 4 ] − . Thiocyanate forms blood-red [Fe(SCN)(H 2 O) 5 ] 2+ , a test for iron(III). [71] Most iron(III) complexes are high-spin, except those with ligands high in the spectrochemical series, like cyanide. [Fe(CN) 6 ] 3− is low-spin. Iron exhibits all possible spin states for d-block elements, from 0 to 5/2. [67]
Iron(II) complexes are less stable but prefer O-donor ligands less. [72]
Organometallic Compounds
[Image of iron pentacarbonyl]
Organoiron chemistry studies iron compounds with covalent carbon bonds. These include cyanide complexes, carbonyl complexes, and sandwich compounds.
[Image of Prussian blue]
Prussian blue, or ferric ferrocyanide, is an old pigment. [63] Iron pentacarbonyl, Fe(CO) 5, yields carbonyl iron powder. [78]
[Image of ferrocene structure and powder]
Ferrocene, discovered in 1951, is a remarkably stable sandwich compound. [79][80][81][82] It remains a key model in the field. [83]
Iron catalysts are used in reactions like transfer hydrogenation. [84]
Industrial Uses
As Structural Material
Iron is the most used metal, over 90% of world production. [129] It's strong and cheap, used in machinery, rails, automobiles, ship hulls, rebar, and buildings. Pure iron is soft, so it's usually alloyed into steel. [130]
Mechanical Properties
[Table of tensile strength and Brinell hardness of iron forms]
Mechanical properties are critical. Pure iron is soft, even softer than aluminum in single-crystal form. [127] Industrial pure iron has a hardness of 20–30 Brinell. [132]
Carbon content significantly increases hardness and strength. [133] Iron is easier to work than ruthenium or osmium. [17]
Types of Steels and Alloys
[Iron-carbon phase diagram]
α-Iron dissolves little carbon (0.021% max at 910 °C). [134] Austenite (γ-iron) dissolves more (2.04% max at 1146 °C) and is used in stainless steel. [21]
Iron is classified by purity and additives. [Pig iron] has high carbon (3.5–4.5%) and contaminants. It's an intermediate for cast iron and steel. [135]
"White" cast irons have carbon as cementite (Fe 3 C), making them hard and brittle. [17] Slow cooling of iron with 0.8% carbon yields pearlite (layers of cementite and α-iron). Rapid cooling creates hard, brittle martensite. [17]
[Gray iron] has carbon as graphite flakes, creating stress concentration points. [Ductile iron] has graphite as spheroids, making it tougher. [136]
[Wrought iron] has low carbon (<0.25%) and slag, giving it a fibrous texture. It's more corrosion-resistant than steel but largely replaced by mild steel. [Carbon steel] has ≤2% carbon. [137] Alloy steels contain other metals, raising costs for specialized uses. Stainless steel is common. High-strength, low-alloy (HSLA) steels offer good properties at low cost. [137][138][139]
High-purity iron alloys have enhanced properties like ductility, tensile strength, and corrosion resistance.
Iron is also used for radiation shielding, being stronger than lead mechanically. [140]
Rusting is a major drawback, costing over 1% of the world economy. [141] Protection methods include painting, galvanization, and passivation. Rusting mechanism: Cathode: 3 O 2 + 6 H 2 O + 12 e − → 12 OH − Anode: 4 Fe → 4 Fe 2+ + 8 e − ; 4 Fe 2+ → 4 Fe 3+ + 4 e − Overall: 4 Fe + 3 O 2 + 6 H 2 O → 4 Fe(OH) 3 or 4 FeO(OH) + 4 H 2 O The electrolyte is often iron(II) sulfate or salt particles. [141]
Catalysts and Reagents
Iron is inexpensive and non-toxic, leading to research in Fe-based catalysts. [142] It's less common as a commercial catalyst than other metals, but pervasive in biological enzymes. [143]
Iron catalysts are used in the Haber–Bosch process and Fischer–Tropsch process. [144] Powdered iron is used in the Bechamp reduction of nitrobenzene to aniline. [145]
Iron Compounds
Iron(III) oxide with aluminium powder creates the thermite reaction, used for welding and ore purification. Iron oxides are used as pigments.
Iron(III) chloride is used in water purification, dyeing, animal feed, and as an etchant for copper in printed circuit boards. [146]
Iron(II) sulfate is a precursor to other iron compounds and used to fortify foods and treat iron deficiency anemia. [146]
Sodium nitroprusside, a vasodilator, is on the World Health Organization's List of Essential Medicines. [148]
Biological and Pathological Role
Iron is essential for life. [10][149][150] Iron–sulfur clusters are vital in enzymes like nitrogenase. Iron proteins handle oxygen transport and electron transfer. [151]
[Heme B structure]
In humans, examples include hemoglobin, cytochrome, and catalase. [10][152] An adult has about 4 grams of iron, mostly in hemoglobin. The body recycles hemoglobin iron. [151][153]
Microbial growth can be affected by iron oxidation/reduction. [154]
Biochemistry
Iron acquisition is challenging because ferric iron is poorly soluble. Organisms use siderophores to bind and transport iron. [155][156][157]
In humans, transferrin binds and transports iron from the duodenum to cells. [10][159] Transferrin's high stability constant ensures efficient iron uptake. In bone marrow, transferrin is reduced, and iron is stored as ferritin for hemoglobin synthesis. [151]
Heme proteins like hemoglobin and myoglobin are key iron compounds. They transport gases, form enzymes, and transfer electrons. [151] Metalloproteins like ferritin and rubredoxin also contain iron. [151] Enzymes like catalase, [160] lipoxygenases, [161] and IRE-BP are vital. [162]
Hemoglobin carries oxygen from lungs to muscles, where it's transferred to myoglobin for storage. [10] Hemoglobin transports CO 2 back to the lungs. [151] In hemoglobin, iron is in a heme group. Its coordination sites bind oxygen reversibly. [151] When oxygen-free (deoxyhemoglobin), Fe 2+ is high-spin, and the heme group bends. [151]
Oxygen binding causes Fe 2+ to switch to low-spin, shrinking the ion and flattening the heme group. This binding also triggers cooperative changes in hemoglobin's subunits, increasing oxygen affinity. Myoglobin lacks this effect. The Bohr effect describes how CO 2 reduces hemoglobin's oxygen affinity. [151]
[Carboxyhemoglobin structure]
Carbon monoxide and phosphorus trifluoride are toxic because they bind hemoglobin more strongly than oxygen. Cyanide mainly interferes with cytochrome a. [151] Cytochromes use heme groups for glucose metabolism. Electron transfer occurs as iron cycles between +2 and +3 states. [151]
Iron–sulfur proteins are crucial for electron transfer, as iron can exist in +2 and +3 states. They contain multiple iron atoms coordinated to sulfur. Rubredoxin is the simplest. Ferredoxins have multiple iron atoms. [151]
Mussels use iron-based bonds in their cuticles for grip. [164]
Nutrition
Diet Rich sources of iron include red meat, oysters, beans, poultry, fish, leaf vegetables, tofu, and blackstrap molasses. [10] Foods like bread and breakfast cereals are often fortified. [10][165]
Dietary supplements often use iron(II) fumarate, though iron(II) sulfate is cheaper and equally absorbed. [146] Elemental iron is also used. Iron is best absorbed when chelated to amino acids, like glycine. [167][168]
Dietary Recommendations RDAs vary by age and sex. For women aged 19-50, it's 18 mg/day. Pregnancy: 27 mg/day. [10] ULs (Tolerable upper intake levels) are set to prevent toxicity. [169] European recommendations (PRIs) are similar but slightly higher. [170]
Infants on cow's milk may need supplements. Frequent blood donors should monitor iron levels. [173]
For U.S. labeling, 100% DV is 18 mg. [174][175]
Deficiency
[Iron deficiency] is the most common nutritional deficiency globally. [10][176][177][178] Untreated deficiency leads to iron-deficiency anemia. Children and pre-menopausal women are most susceptible. Severe cases can cause heart problems, pregnancy complications, and developmental delays. [180]
The brain is somewhat protected from acute deficiency. [181] Prolonged deficiency may reduce brain iron, impacting neurotransmitter synthesis and myelin. [184] Animal studies show changes resembling Parkinson's and Huntington's. [185][186] However, iron accumulation in the brain is also linked to Parkinson's. [187]
Excess
[Iron overload] occurs when the body can't regulate iron uptake properly, often due to genetic defects affecting hepcidin. [189] This can lead to hemochromatosis. [10] People should not take iron supplements without medical advice. [190]
Excess iron creates reactive free radicals, damaging cells. Toxicity occurs when iron exceeds transferrin's binding capacity. Organs like the heart and liver can be damaged. [191] Lethal dose is around 60 mg/kg. [192] Iron poisoning is a common cause of death in young children. [192] DRI sets the UL at 45 mg/day for adults. [193]
Deferoxamine is a chelating agent used to treat iron toxicity. [194][195]
ADHD
Some research links low thalamic iron to ADHD. [196] Iron supplementation may help, especially the inattentive subtype. [197] A 2012 study found no correlation between serum ferritin and ADHD. [198]
Cancer
Iron's role in cancer is complex. Chemotherapy can cause iron deficiency, treated with IV iron. [200] Iron overload, possibly from red meat consumption, may promote tumor growth and cancer susceptibility, particularly colorectal cancer. [10][200]
Marine Systems
Iron is a limiting nutrient for plankton in oceans. [201] Too little iron can slow phytoplankton growth. [202] Iron can enter oceans via rivers and the atmosphere. [204] Sea ice in the Arctic stores and releases iron. [205] The iron cycle fluctuates between aqueous and particle forms. [206]