Iron

What is Iron?

Elemental iron possesses the atomic number 26 and the symbol Fe, derived from the Latin ferrum, meaning “iron.”

The metal in question is a member of group 8 and the first transition series of the periodic table. It makes up a large portion of the outer and inner core of the planet and is, by mass.

The most abundant element on Earth, slightly more so than oxygen (32.1% and 30.1%, respectively).

As the fourth most prevalent element in the crust of the Earth, it is primarily deposited in its metallic state by meteorites, although its ores are also found there.

Kilns or furnaces that can reach 1,500 °C (2,730 °F) or more—roughly 500 °C (932 °F) greater than what is needed to smelt copper—are necessary to extract useable metal from iron ores.

In Eurasia, humans began to master that technique in the second millennium BC, and copper alloys started to lose ground to iron-based tools and weapons—in some areas, as late as 1200 BC.

That incident is seen as the shift. from the Age of Bronze to the Age of Iron. Because of their affordable price and mechanical qualities.

Iron alloys, including cast iron, stainless steel, steel, and special steels, are by far the most widely used industrial metals in the modern world.

Due to its low cost—just a few dollars per kilogram or pound—iron is the most affordable metal, making the iron and steel sector highly significant economically.

The smooth, pristine surfaces of pure iron have a silvery-gray reflection. Rust is the result of the easy reaction of iron with oxygen and water to form brown-to-black hydrated iron oxides.

Rust occupies a larger volume than the metal it flakes off, exposing more surfaces to corrode than the oxides of some other metals that form passivating layers.

Higher purity irons, like electrolytic iron, have greater corrosion resistance. An adult human’s body contains around 4 grams (0.005% body weight) of iron, primarily in the form of hemoglobin and myoglobin.

Since these two proteins carry oxygen through the blood and store it in the muscles, respectively, they are essential for vertebrates’ metabolism.

Human iron metabolism requires a minimum amount of iron in the diet to sustain the required levels.

Many significant redox enzymes involved in cellular respiration, as well as oxidation and reduction in plants and animals, have iron as their metal in their active site.

Iron(II) and iron(III) are the two most prevalent oxidation states of iron chemically. Many characteristics of iron are also shared by other transition metals, such as ruthenium and osmium, two other elements in group 8.

From -4 to +7, iron can form compounds in a wide variety of oxidation states. Numerous coordination compounds are also formed by iron.

These include ferrocene, ferrioxalate, and Prussian blue, all of which have important uses in industry, medicine, or research.

The most notable element that might be produced through exothermic nucleosynthesis is iron, which has the highest atomic number.

Characteristics

Allotropes

Different atom configurations inside the solid form at least four known allotropes of iron, which are commonly identified as a,? d, and e.

At standard pressures, the first three types are seen. Iron crystallizes into its d allotrope, which has a body-centered cubic (bcc) crystal structure when it cools below its freezing point of 1538 °C.

It transitions to its ?-iron allotrope, or austenite, a face-centered cubic (fcc) crystal structure, as it cools down to 1394 °C.

The crystal structure reverts to the bcc a-iron allotrope at temperatures lower than 912 °C. Due to their importance to ideas regarding the composition of the.

Earth’s and other planets’ cores, the physical characteristics of iron under extreme pressure and temperature conditions have also been thoroughly investigated.

At temperatures of a few hundred kelvin or below, and at around 10 GPa, a-iron transforms into another hexagonal close-packed (hcp) structure called e-iron.

Though at a higher pressure, the ?-phase, which is at a higher temperature, also transforms into an e-iron.

There is debatable experimental support for a stable ß phase at pressures more than 50 GPa and temperatures greater than 1500 K.

Its structure should be either double hcp or orthorhombic. Confoundingly, even though a-iron’s crystal structure remains unchanged, it is frequently referred to as “ß-iron” above the Curie point, when it transitions from ferromagnetic to paramagnetic.

It is widely assumed that an iron-nickel alloy with an e (or ß) structure makes up the Earth’s inner core.

Melting and Boiling Points

Iron’s melting and boiling points, as well as its enthalpy of atomization, are lower than those of the preceding 3d elements, ranging from scandium to chromium.

This indicates that the 3d electrons’ contribution to metallic bonding is diminished as they are drawn further into the inert core by the nucleus; on the other hand.

They are higher than those of the preceding element, manganese because that element has a partially filled 3d sub-shell, which makes its d-electrons difficult to delocalize.

For ruthenium, but not for osmium, a similar pattern is seen. For pressures below 50 GPa, iron’s melting point is clearly known by experiment.

Published data (as of 2007) still differs by tens of gigapascals and more than a thousand kelvin for higher pressures.

Magnetic Properties

a-iron becomes ferromagnetic below its Curie point of 770 °C (1,420 °F; 1,040 K) when the two unpaired electrons in each atom create an overall magnetic field by generally aligning with their neighbors’ spins.

This occurs because those two electrons’ orbitals (d_z2 and dx² -. y²) do not point in the direction of nearby atoms in the lattice and are hence not engaged in metallic bonding.

The atoms divide spontaneously into magnetic domains, each measuring roughly 10 micrometers across, with atoms in each domain having parallel spins, while some domains have different orientations in the absence of an external magnetic field source.

Thus, the total magnetic field of a macroscopic piece of iron will be almost zero. when a magnetic field is applied externally.

Adjacent domains pointing in different directions are forced to expand at the expense of those that are magnetized in the same general direction, thus intensifying the external field.

This effect is used in electrical transformers, magnetic recording heads, and electric motors, among other devices that must channel magnetic fields in order to perform their intended functions.

The iron item can become a (permanent) magnet when impurities, lattice defects, or grain and particle boundaries “pin” the domains in their new locations.

This effect lasts long after the external field is eliminated. Certain iron compounds, such as ferrites and the mineral magnetite.

Which is a crystalline form of the mixed iron(II, III) oxide Fe₃O₄, displays similar characteristics (though ferrimagnetism, the atomic-scale process, is slightly different).

The first navigational compasses were made of magnetite fragments with naturally occurring permanent magnetization, or lodestones.

Before cobalt-based materials took their place, magnetite particles were widely employed in magnetic recording media such as disks, floppies, magnetic tapes, and core memory.

Isotopes

There are four stable isotopes of iron: ⁵⁷Fe (2.119%), ⁵⁶Fe (91.754%), ⁵⁸Fe (0.282%), and ⁵⁴Fe (5.845% of natural iron).

Additionally, twenty-four synthetic isotopes have been produced. ⁵⁷Fe is the sole stable isotope with a nuclear spin of –¹/₂.

Though the process of double electron capture from the nuclide ⁵⁴Fe to ⁵⁴Cr has never been observed, only a lower limit of 3.1×1022 years has been established for the half-life.

⁶⁰Fe is a long-half-life (2.6 million years) radionuclide that is now extinct. Although it is not present on Earth, its granddaughter, the stable nuclide ⁶⁰Ni, is the final byproduct of its decay.

The nucleosynthesis of ⁶⁰Fe through studies of meteorites and ore formation has dominated much of the previous research on the isotopic composition of iron.

Mass spectrometry developments during the past ten years have made it possible to identify and measure minute, naturally occurring differences in the ratios of the stable isotopes of iron.

Applications to biological and industrial systems are beginning to emerge, although the Earth and planetary science groups still drive a large portion of this work.

Evidence for the presence of ⁶⁰Fe during the Solar System’s creation was found in the phases of the meteorites.

Chervony Kut and Semarkona, where a correlation was observed between the stable iron isotope abundance and the concentration of ⁶⁰Ni, which is ⁶⁰Fe’s granddaughter.

It’s possible that once asteroids formed 4.6 billion years ago, the energy supplied by the disintegration of ⁶⁰Fe and ²⁶Al helped with the remelting and differentiation of asteroids.

The abundance of ⁶⁰Ni found in extraterrestrial material could provide more information about the Solar System’s early history and origin.

Since it is the most frequent endpoint of nucleosynthesis, nuclear scientists are especially interested in the most prevalent iron isotope, ⁵⁶Fe.

It is the endpoint of fusion chains inside extremely massive stars because ⁵⁶Ni (14 alpha particles) is easily produced from lighter nuclei in the alpha process in nuclear reactions in supernovae (see silicon burning process).

The addition of another alpha particle, resulting in ⁶⁰Zn, requires a great deal more energy. These stars produce large amounts of this ⁵⁶Ni, which has a half-life of roughly six days, but it decays by two.

In the supernova remnant gas cloud, there are sequential positron emissions from supernova decay products that first produce radioactive ⁵⁶Co and subsequently stable ⁵⁶Fe.

Iron is therefore the most common element in red giants’ cores, as well as the most common metal in iron meteorites and the thick metal cores of planets like Earth.

Comparatively speaking to other stable metals with roughly the same atomic weight, it is also quite frequent in the cosmos.

Iron is the most common refractory element and the sixth most plentiful element in the universe.

Star circumstances are not favorable for the synthesis of ⁶²Ni, despite the fact that this material has slightly greater binding energy than ⁵⁶Fe and may yield a further small energy gain.

Iron predominates over nickel in supernova element synthesis, yet ⁵⁶Fe nonetheless has a lower mass per nucleon than ⁶²Ni because of its higher proportion of lighter protons.

Therefore, elements heavier than iron can only originate during a supernova, which involves 56Fe nuclei rapidly capturing neutrons.

Cold fusion via quantum tunneling would allow the light nuclei in the ordinary matter to fuse into ⁵⁶Fe nuclei in the remote future of the universe, providing that proton decay does not occur.

All stellar-mass objects would then degrade into iron by fission and alpha particle emission, resulting in frigid spheres composed entirely of iron.

Origin and Occurrence in Nature

Cosmogenesis

Because iron is produced in large quantities during the uncontrolled fusion and explosion of type Ia supernovae, which disperse iron throughout space, rocky planets like Earth have an abundance of iron.

Metallic Iron

Because it oxidizes easily, metallic or native iron is rarely found on Earth’s surface. Nonetheless, it is thought that the majority of the material in the.

Earth’s inner and outer cores, which combined make up 35% of the planet’s mass, are an iron alloy, maybe containing nickel.

The Earth’s magnetic field is thought to have originated from electric currents in the liquid outer core.

It is thought that the Moon and the four terrestrial planets, Mercury, Venus, and Mars, have metallic cores made primarily of iron.

It is also thought that metallic iron alloy makes up the majority or portion of the M-type asteroids.

The primary source of naturally occurring metallic iron on Earth’s surface is the rare iron meteorite.

It has been reported that the Inuit in Greenland used iron from the Cape York meteorite for tools and hunting weapons.

Artifacts composed of cold-worked meteoritic iron have been discovered in a number of archeological sites that date from a period when iron smelting had been invented.

One in twenty meteorites are thought to contain the special iron-nickel minerals. Kamacite has 90–95% iron, while taenite has 35–80% iron.

Additionally, native iron is rarely found in basalts generated from magmas that have come into contact with sedimentary rocks rich in carbon, as the carbon-rich rocks have sufficiently lowered the oxygen fugacity to allow the iron to crystallize.

This is referred to as telluric iron and has been reported from a few locations, including Bühl, Germany, Yakutia, Russia, and Disko Island, West Greenland.

Mantle Minerals

Mg, Fe ferropericlaseAbout 20% of the Earth’s lower mantle is composed of O, a solid solution of periclase (MgO) and wüstite (FeO).

This makes it the second most abundant mineral phase in the lower mantle after silicate perovskite (Mg, Fe)SiO₃.

O is also the main host for iron in the lower mantle..γ-(Mg, Fe)₂[SiO₄] ↔ (Mg, Fe)[SiO₃] + (Mg, Fe)O  at the base of the mantle’s transition zone converts ?-olivine into a combination of ferropericlase with silicate perovskite, and vice versa.

This lower mantle mineral phase is also frequently referred to as magnesiowüstite in the literature.

Up to 93% of the lower mantle may be composed of silicate perovskite, and 38% of the Earth’s volume is made up of the magnesium iron form, or (Mg, Fe)SiO₃, which is thought to be the most common mineral.

Earth’s Crust

The majority of the iron on Earth is found in the inner and outer cores, despite being the most abundant element overall.

Although iron makes up only 5% of the crust’s total mass, it is the fourth most abundant element in that layer of the Earth’s crust (after oxygen, silicon, and aluminum).

Numerous iron minerals are formed when the majority of the iron in the crust combines with different elements.

The iron oxide minerals, which include siderite (FeCO₃), magnetite (Fe3O₄), and hematite (Fe₂O₃), are a significant class of minerals.

Pyrrhotite and pentlandite are two more sulfide minerals found in many igneous rocks. Iron tends to leach during weathering as the sulfate from sulfide deposits and as the bicarbonate from silicate deposits.

In an aqueous solution, both of these oxidize and precipitate as iron(III) oxide, even at slightly raised pH levels.

Banded iron formations, a type of rock with repeated thin layers of iron oxides alternating with bands of iron-poor shale and chert, are where large iron deposits are found.

The period between 3,700 and 1,800 million years ago is when the banded iron formations were formed.

Since prehistoric times, materials containing finely crushed iron(III) oxides or oxide-hydroxides, like ochre, have been employed as pigments in the colors yellow, red, and brown.

They also influence the color of different rocks and clays, including whole geological formations such as the Buntsandstein (“colored sandstone”, British Bunter) and the Painted Hills in Oregon.

The yellowish hue of numerous historical buildings and sculptures can be attributed to iron compounds, such as those found in Bathstone in the UK and Eisensandstein, a Jurassic “iron sandstone” that comes from places like Donzdorf in Germany.

Mars’s surface is famously red because of regolith which is rich in iron oxide. The iron sulfide mineral pyrite (FeS₂) contains significant amounts of iron, but it is not used since it is difficult to extract.

Iron is so abundant in fact that ores with extremely high iron content are usually the only ones used in production.

The Metal Stocks in Society report by the International Resource Panel states that there are 2,200 kg of iron in use worldwide per person.

In this regard, more developed nations differ from less developed nations (7,000–14,000 vs 2,000 kg per capita).

Oceans

Ocean science has shown how iron influenced the temperature and marine biota of the ancient oceans.

Chemistry and Compounds

Iron exhibits the typical chemical characteristics of transition metals, including large coordination, and the ability to form variable oxidation states with differences of only one step.

And organometallic chemistry—the latter field was revolutionized in the 1950s by the discovery of an iron compound called ferrocene.

Iron’s availability and significant contribution to human technological advancement have led to the idea that iron represents a prototype for the entire transition metal block.

Its 26 electrons are grouped in the configuration [Ar]3d⁶4s², where a number of electrons can be ionized due to the relative energy of the 3d and 4s electrons.

Oxidation state	Representative compound
−4 (d10s2)	[FeIn6−xSnx]
−2 (d10)	Disodium tetracarbonylferrate (Collman’s reagent)
−1 (d9)	Fe2(CO)2−8
0 (d8)	Iron pentacarbonyl
1 (d7)	Cyclopentadienyliron dicarbonyl dimer (“Fp2“)
2 (d6)	Ferrous sulfate, Ferrocene
3 (d5)	Ferric chloride, Ferrocenium tetrafluoroborate
4 (d4)	Fe(diars)2Cl2+2, FeO(BF4)2
5 (d3)	FeO3−4
6 (d2)	Potassium ferrate
7 (d1)	[FeO4]– (matrix isolation, 4K)

The two main oxidation states in which iron forms compounds are +2 (iron(II), “ferrous”) and +3 (iron(III), “ferric”).

Additionally, iron can be found in higher oxidation states. For example, purple potassium ferrate (K₂FeO₄) includes iron in the +6 oxidation state.

After laser-ablated Fe atoms were condensed with a combination of O2/Ar, the anion [FeO₄]–containing iron in its +7 oxidation state and an iron(V)-peroxo isomer were discovered by infrared spectroscopy at 4K.

Often used as an intermediary in several oxidation reactions in biology. +1, 0, -1, or even -2 formal oxidation states are present in many organic compounds.

The method known as Mössbauer spectroscopy is frequently employed to evaluate the oxidation states and additional bonding characteristics.

Iron(II) and iron(III) centers can be found in many mixed-valence compounds, such as magnetite and Prussian blue (Fe₄(Fe[CN]₆)₃).

The latter serves as the conventional “blue” in architectural drawings. While its heavier congeners, ruthenium, and osmium, can achieve their group oxidation state of +8, iron is the first transition metal that is unable to do so, with ruthenium having greater trouble than osmium.

Osmium prefers high oxidation states where it forms anionic complexes, whereas ruthenium exhibits an aqueous cationic chemistry in its low oxidation states that is similar to that of iron.

The vertical similarities down the groups in the second half of the 3d transition series rival the horizontal similarities between iron and its neighbors, cobalt and nickel, in the periodic table, which are likewise ferromagnetic at room temperature and have comparable chemistry.

For this reason, the iron triad—iron, cobalt, and nickel—is occasionally used interchangeably. Iron and mercury do not combine to produce amalgams, in contrast to many other metals.

Mercury is therefore traded in iron-based, standardized 76-pound (34-kg) flasks. Among the elements in its group, iron is by far the most reactive; when divided finely, it becomes pyrophoric and dissolves readily in dilute acids to form Fe²⁺.

Because an impermeable oxide layer forms, it cannot react with concentrated nitric acid or other oxidizing acids, but it can react with hydrochloric acid. Because of its oxide coating, high-purity iron, also known as electrolytic iron, is said to be rust-resistant.

Binary Compounds

Oxides and sulfides

The most prevalent iron(II, III) oxide (Fe₃O₄) and iron(III) oxide (Fe₂O₃) are among the many oxide and hydroxide compounds that iron can produce.

There is also iron(II) oxide, however it is unstable at ambient temperature. They are all non-stoichiometric compounds with variable compositions, despite their names.

The main ores used to produce iron are these oxides (see bloomery and blast furnace). They are also utilized in the synthesis of pigments and ferrites, which are helpful magnetic storage devices for computers.

Due to its golden sheen, iron pyrite (FeS₂), commonly referred to as fool’s gold, is the most well-known sulfide.

It is actually an iron(II) polysulfide with Fe²⁺ and S₂^-2 ions in a deformed sodium chloride structure rather than an iron(IV) molecule.

Halides

Ferrous and ferric halides are known to exist in binary form. Usually, iron metal is treated with the appropriate hydrohalic acid to produce the matching hydrated salts, which is how the ferrous halides are produced.

Fe + 2 HX → FeX₂ + H₂ (X = F, Cl, Br, I)

Ferric halides, of which ferric chloride is the most prevalent, are produced when iron combines with fluorine, chlorine, and bromine.

2 Fe + 3 X₂ → 2 FeX₃ (X = F, Cl, Br)

A notable exception is ferric iodide, which is thermodynamically unstable because of the strong oxidizing power of Fe3+ and the strong reducing power of I-:

2 I⁻ + 2 Fe³⁺ → I₂ + 2 Fe²⁺ (E⁰ = +0.23 V)

The black solid known as ferric iodide is not stable under normal circumstances. However, it can be made by reacting iron pentacarbonyl with iodine and carbon monoxide in.

The presence of hexane, light, and a temperature of -20 °C without the presence of oxygen or water. Ferric iodide complexes with certain soft bases are recognized to be stable substances.

Solution Chemistry

The following lists the standard reduction potentials for a few typical iron ions in an acidic aqueous solution:

[Fe(H2O)6]2+ + 2 e−	⇌ Fe	E0 = −0.447 V
[Fe(H2O)6]3+ + e−	⇌ [Fe(H2O)6]2+	E0 = +0.77 V
FeO2− 4 + 8 H3O+ + 3 e−	⇌ [Fe(H2O)6]3+ + 6 H2O	E0 = +2.20 V

Because of its extreme oxidizing power, the reddish-purple tetrahedral ferrate(VI) anion oxidizes both water and ammonia to nitrogen (N₂).

4 FeO²⁻₄ + 34 H₂O → 4 [Fe(H₂O)₆]³⁺ + 20 OH⁻+ 3 O₂

The hexaquo complex pale-violet [Fe(H₂O)₆]³⁺ acid with a value is fully hydrolyzed above pH 0:

[Fe(H₂O)₆]³⁺⇌ [Fe(H₂O)₅(OH)]²⁺ + H⁺K = 10^−3.05 mol dm⁻³[Fe(H₂O)₅(OH)]²⁺⇌ [Fe(H₂O)₄(OH)₂]⁺ + H⁺K = 10^−3.26 mol dm⁻³2[Fe(H₂O)₆]³⁺⇌ [Fe(H₂O)₄(OH)]4+2 + 2H⁺ + 2H₂OK = 10^−2.91 mol dm⁻³

The above yellow hydrolyzed species arise when pH rises over 0 and reddish-brown hydrous iron(III) oxide precipitates out of solution when pH rises above 2-3.

Fe³⁺ has a d⁵ configuration, but because of its higher positive charge and greater polarization, which lowers the energy of its ligand-to-metal charge transfer absorptions, Fe³⁺ has a different absorption spectrum than Mn²⁺, which contains weak, spin-forbidden d–d bands.

With the exception of the hexaquo ion, which has a spectrum dominated by charge transfer in the near ultraviolet range, all of the aforementioned complexes have rather vivid colors.

However, there is no significant hydrolysis of the pale green iron(II) hexaquo ion [Fe(H₂O)₆]²⁺. When carbonate anions are introduced, white iron(II) carbonate precipitates out instead of carbon dioxide evolving.

When carbon dioxide is present in excess, this produces somewhat soluble bicarbonate, which is frequently found in groundwater.

However, it oxidizes fast in the air to make iron(III) oxide, which is responsible for the brown deposits found in a significant number of streams.

Coordination Compounds

The electronic structure of iron results in a very extensive coordination and organometallic chemistry.

Geometric isomers exist in complexes containing several bidentate ligands. For instance, the trans-chlorohydridobis(bis-1,2-(diphenylphosphine)ethane)iron(II) complex is employed as a precursor for Fe(dppe)₂ moiety-containing compounds.

Three oxalate ligands make up the ferrioxalate ion. (shown at right) exhibits helical chirality with its two non-superposable geometries labeled, according to IUPAC norms, G (delta) for the right-handed screw axis and ? (lambda) for the left-handed screw axis.

In chemical actinometry, potassium ferrioxalate is utilized, and in traditional photographic procedures, it is subjected to photoreduction alongside its sodium salt.

The iron(II) oxalate dihydrate, as shown below, has a polymeric structure with co-planar oxalate ions bridging between iron centers and the water of crystallization forming the caps of each octahedron.

The only difference between iron(III) and chromium(III) complexes is that iron(III) prefers O-donor ligands over N-donor ligands.

The latter frequently dissociate in water and have a tendency to be somewhat more unstable than iron(II) complexes.

Numerous Fe–O compounds exhibit vivid colors and are employed in phenol and enol testing.

For instance, iron(III) chloride and phenol combine to generate a deep violet complex in the ferric chloride test, which is used to detect the presence of phenols:
3 ArOH + FeCl₃ → Fe(OAr)₃ + 3 HCl (Ar = aryl)

There are numerous iron coordination compounds known. Hexachloroferrate(III), [FeCl₆]³⁻, is a common six-coordinate anion that is present in the mixed salt tetrakis(methylammonium) hexachloroferrate(III) chloride.

Fluoro complexes of iron(III) are the most stable of the halide and pseudohalide complexes; in aqueous solution, the colorless [FeF₅(H₂O)]²⁻ is the most stable.

Chloro complexes, such as [FeCl₄]⁻;  [FeBr₄]⁻, and [FeI₄]⁻, are less stable and prefer tetrahedral coordination because they can be readily reduced to iron(II).

Since thiocyanate creates the blood-red [Fe(SCN)(H₂O)₅]²⁺, it is a commonly used test for the presence of iron(III).

The majority of iron(III) complexes, like manganese(II), are high-spin, with the exception of compounds like cyanide that have ligands at the top of the spectrochemical series. [Fe(CN)₆]³⁻ is an illustration of a low-spin iron(III) complex.

The electronic spin states of iron are extremely diverse, exhibiting all possible spin quantum numbers for d-block elements, ranging from 0 (diamagnetic) to  ⁵⁄₂ (5 unpaired electrons).

There are always half as many unpaired electrons as this value. Low-spin complexes have zero to two unpaired electrons, whereas high-spin complexes have four or five.

Although O-donor ligands are preferred less strongly in iron(II) complexes than in iron(III) complexes, iron(II) complexes are less stable than iron(III) complexes; for instance, [Fe(NH₃)₆]²⁺ is known whereas [Fe(NH₃)₆]³⁺ is not.

Although they have a propensity to oxidize to iron(III), low pH and the particular ligands employed can mitigate this.

Organometallic Compounds

The study of iron compounds known as organometallics, in which the carbon atom is covalently bonded to the metal atom, is known as organoiron chemistry.

They come in a wide variety of forms, such as sandwich and half-sandwich compounds, carbonyl complexes, and cyanide complexes.

 Fe₄[Fe(CN)₆]₃ often known as “ferric ferrocyanide” or Prussian blue, is a well-known and ancient iron-cyanide complex that is widely utilized as a pigment and in a variety of other applications.

Its creation can be utilized as a straightforward wet chemistry test to differentiate between Fe²⁺ and Fe³⁺ water solutions as they react with potassium ferricyanide and potassium ferrocyanide, respectively, to generate Prussian blue.

Iron pentacarbonyl, or Fe(CO)₅, is another older example of an organoiron chemical. It is a compound in which five carbon monoxide molecules’ carbon atoms are linked to a neutral iron atom.

The substance can be utilized to create carbonyl iron powder, a very reactive metallic iron form. Triiron dodecacarbonyl,  Fe₃(CO)₁₂, a complex with a core consisting of three iron atoms, is produced by the thermolysis of iron pentacarbonyl.

Disodium tetracarbonylferrate, often known as Collman’s reagent, is a helpful tool for organic chemistry since it includes iron in the -2 oxidation state.

Iron is present in the uncommon+1 oxidation state in cyclopentadienyl iron dicarbonyl dimer. The unusually stable sandwich compound ferrocene Fe(C₅H₅)₂, discovered in 1951 by Pauson and Kealy and independently by Miller and others, was a landmark in this research.

Woodward, Wilkinson, and Fischer confirmed the compound’s astonishing molecular structure barely a year later.

One of the most crucial resources and models in this class is still ferrocene. Catalysts comprise iron-centered organometallic substances. One catalyst for the transfer hydrogenation of ketones is the Knölker complex.

Industrial uses

In the industrial setting, iron(II) sulfate (FeSO₄·7H₂O) and iron(III) chloride (FeCl₃) are the most widely produced iron compounds.

While Mohr’s salt ((NH₄)₂Fe(SO₄)₂·6H₂O). is more stable to aerial oxidation, the former is one of the most accessible sources of iron(II).

In the air, iron(II) molecules often oxidize to iron(III) compounds.

History

Development of Iron Metallurgy

Undoubtedly, one of the elements that the ancient world was aware of was iron. For millennia, people have worked, or wrought, it.

Unfortunately, because iron corrodes easily, antique iron artifacts are far rarer than gold or silver ones.

The advancement of technology was gradual, and it took several centuries for iron to overtake bronze as the preferred metal for tools and weapons, even with the discovery of smelting.

Meteoritic Iron

G.A. Wainwright discovered beads made of meteoric iron in Gerzeh, Egypt, dating back to 3500 BC or before.

The beads’ 7.5% nickel content is indicative of their meteoric origin, as iron present in the Earth’s crust typically only contains trace amounts of nickel impurities.

Because it originated in the sky, meteoric iron was highly valued and frequently used to create tools and weapons.

For instance, a meteoric iron knife recovered in Tutankhamun’s tomb had the same amounts of iron, cobalt, and nickel as a meteorite found nearby that had been left behind by a prehistoric meteor shower. Artifacts dated 3000 to 2500 BC are probably Egyptian creations wrought from iron.

Because of the nickel component, meteoritic iron is relatively soft, ductile, and easily cold forged, but when heated, it may become brittle.

Wrought Iron

Although iron was initially produced during the Middle Bronze Age, it took several centuries for iron to completely replace bronze.

Between 3000 and 2700 BC, samples of smelted iron were made in Tall Chagar Bazaar in northern Syria and Asmar, Mesopotamia. About 1600 BC, the Hittites founded an empire in north-central Anatolia.

They seem to have been the first in their civilization to recognize the value of iron and learn how to extract it from its ores.

Between 1500 and 1200 BC, the Hittites started to smelt iron, and after their empire collapsed in 1180 BC, the practice spread to the remainder of the Near East.

The Iron Age is the term for the era that followed. Smelted iron artifacts from 1800–1200 BC have been discovered in India, and artifacts from approximately 1500 BC have been discovered in the Levant, indicating smelting in Anatolia or the Caucasus.

Claims that iron was utilized in India very early have been supported by allusions to iron in the Indian Vedas, which have been used to date the books themselves (see History of Metallurgy in South Asia).

The term “metal” in the Rigveda refers to copper, while the post-Rigvedic Atharvaveda is the first to mention iron, which is known as “black copper” or syama ayas.

According to certain archeological findings, iron was melted as early as the seventh century BC in Zimbabwe and southeast Africa.

In the latter part of the 11th century BC, ironworking was brought to Greece, and from there it rapidly expanded across Europe.

Celtic expansion is linked to the growth of ironworking in Central and Western Europe. Iron was widely used during the Roman era, according to Pliny the Elder.

Iron first appeared in the regions that are now known as China between 700 and 500 BC. It’s possible that iron smelting was brought to China from Central Asia.

Cupola furnaces were in use as early as the Warring States period (403–221 BC), and the first blast furnace evidence comes from China in the first century AD.

Throughout the Tang and Song dynasties, blast and cupola furnaces were still widely used. Henry Cort used novel production techniques to start refining iron during the British Industrial Revolution, moving from pig iron to wrought iron (also known as bar iron).

He received a patent for the puddling method of iron ore refinement in 1783. Later, others, notably Joseph Hall, made improvements to it.

Cast Iron

The first cast iron was made in China in the fifth century BC, and it wasn’t until the Middle Ages that cast iron was introduced to Europe.

In what is now Luhe County, Jiangsu, China, researchers found the oldest cast iron objects. Ancient China employed cast iron for construction, agriculture, and warfare.

In Europe during the Middle Ages, methods for turning cast iron—referred to here as “pig iron”—into wrought iron were discovered through the use of finery forges.

Charcoal was necessary as a fuel for each of these procedures. Built of fireproof brick, medieval blast furnaces stood roughly ten feet (3.0 meters) tall.

Hand-operated bellows were typically used to supply forced air. Although they can now generate thousands of tons of iron per day thanks to their larger.

Fourteen-meter-diameter hearths, modern blast furnaces still function largely in the same manner as they did in the Middle Ages.

Although he continued to employ blast furnaces, Abraham Darby I built a coke-fired blast furnace in 1709 to make cast iron, replacing charcoal.

One of the things that led to the Industrial Revolution was the consequent availability of cheap iron.

Due to its lower cost, cast iron started to supplant wrought iron for some uses towards the end of the 18th century.

It wasn’t until the 18th century that the carbon content of iron was suggested as the cause of the variations in wrought iron, cast iron, and steel characteristics.

After the first iron bridge was built in 1778, iron became a prominent structural material as it grew more affordable and more readily available.

As a reminder of the significance iron had throughout the Industrial Revolution, this bridge is still standing today.

After that, iron was utilized for steam engine iron cylinders as well as for rails, boats, ships, aqueducts, and buildings.

A number of languages refer to railways as the “iron road,” including French chemin de fer, German Eisenbahn, Turkish demiryolu, Russian, Chinese, Japanese, Korean ??, and Vietnamese du? ng s?t. Railways have been essential to the construction of modernity and concepts of progress.

Steel

The first steel was made in antiquity using a bloomery; it had a higher carbon content than wrought iron but less than pig iron.

Good steel was being produced by Luristan’s blacksmiths in western Persia by the year 1000 BC. Subsequently, enhanced variations, known as Wootz steel from India and Damascus steel, were created approximately 300 BC and 500 AD, respectively.

Due to the specialized nature of these processes, steel did not gain widespread use until the 1850s.

In the seventeenth century, new techniques for making it were developed by carburizing iron bars during the cementation process.

During the Industrial Revolution, new techniques for making bar iron without the use of charcoal were developed; these techniques were then used to make steel.

Henry Bessemer developed a novel method of producing mild steel in the late 1850s by forcing air through molten pig iron.

Because of this, steel became considerably more affordable, which stopped the mass production of wrought iron.

Foundations of Modern Chemistry

In his experiments that led to the proof of the conservation of mass in 1774, Antoine Lavoisier employed the reaction of water steam with metallic iron inside an incandescent iron tube to produce hydrogen.

This was a crucial step in transforming chemistry from a qualitative to a quantitative science.

Symbolic Role

Iron has a specific place in mythology and is frequently used in folklore and as a metaphor. In order to account for the various eras of humanity.

The Greek poet Hesiod names the various human ages after metals such as gold, silver, bronze, and iron in Works and Days (lines 109–201). Rome and the Iron Age were intimately associated, especially in Ovid’s Metamorphoses

In desperation, the Virtues leave the planet, and human wickedness spreads throughout the entire world. Then, hard steel was successful.

The significance of iron’s symbolic role may be seen in the 1813 German Campaign. At that time, Frederick William III ordered the first Iron Cross to be used as a military medal.

Between 1813 and 1815, when the Prussian royal family encouraged people to donate gold and silver jewelry for military funding.

Berlin iron jewelry manufacture peaked. Later war efforts also made use of the inscription Ich gab money für Eisen, which means “I gave gold for iron.”

Production of Metallic Iron

Laboratory Routes

Pure iron is made in small amounts in the laboratory for specific uses by reducing pure oxide or hydroxide with hydrogen or by producing iron.

Pentacarbonyl and heating it to 250 °C, which causes it to disintegrate and generate pure iron powder. Electrolysis of ferrous chloride onto an iron cathode is an additional technique.

Main Industrial Route

These days, there are two primary steps in the industrial manufacture of iron or steel. The first step involves reducing iron ore and coke in a blast furnace, then separating the molten metal from large contaminants such as silicate particles.

Pig iron, the alloy produced at this step, has a comparatively high carbon content. In the second stage, oxidation reduces the carbon content of the pig iron to produce cast iron, steel, or wrought iron. At this point, alloy steels can be created by adding other metals.

Country	Iron ore	Pig iron	Direct iron	Steel
China	1,114.9	549.4		573.6
Australia	393.9	4.4		5.2
Brazil	305.0	25.1	0.011	26.5
Japan		66.9		87.5
India	257.4	38.2	23.4	63.5
Russia	92.1	43.9	4.7	60.0
Ukraine	65.8	25.7		29.9
South Korea	0.1	27.3		48.6
Germany	0.4	20.1	0.38	32.7
World	1,594.9	914.0	64.5	1,232.4

Blast Furnace Processing

Iron ores, often hematite Fe₂O₃ or magnetite Fe₃O₄, coke (coal that has been individually baked to remove volatile components), and flux (limestone or dolomite) are put into the blast furnace.

The mixture is exposed to “blasts” of air that have been preheated to 900 °C (and occasionally with oxygen enrichment), enough to convert the carbon to carbon monoxide:

{2 C + O2 -> 2 CO}

The temperature rises to roughly 2000 °C as a result of this process. Metallic iron is produced from iron ore by carbon monoxide.
{Fe₂O₃ + 3 CO -> 2 Fe + 3 CO₂}
Coke and some iron in the furnace’s lower, high-temperature area react directly:
{2Fe₂O₃ + 3C -> 4Fe + 3CO₂}
Silica-containing minerals in the ore are eliminated by the flux; they would otherwise jam the furnace.

The carbonates are broken down by the furnace’s heat into calcium oxide, which then combines with any extra silica to create a slag that can contain calcium silicate (CaSiO₃) or other byproducts.

Both the metal and the slag are molten at the temperature inside the furnace. With the slag on top, they gather at the bottom as two immiscible liquid layers that are readily separated.

Slag can be utilized to improve mineral-poor agricultural soils or as a building material for roads.
Thus, one of the industries that contributes the most to global CO₂ emissions is still the steel industry.

Steelmaking

Up to 4-5 percent carbon (by mass) can be found in the pig iron made in a blast furnace, along with trace levels of other impurities like sulfur, magnesium, phosphorus, and manganese.

Because of the high carbon content, it is brittle and rather weak. Steel is created when carbon is reduced to between 0.002 and 2.1%.

Steel has the potential to be up to 1000 times tougher than pure iron. Then, a wide range of steel objects can be produced by machining, hot rolling, cold working, forging, etc.

Cast iron is produced by purging pig iron of impurities while maintaining a 2-4% carbon content. Foundries then use this iron to cast it into products like lamps, rails, radiators, stoves, and pipes.

After being forged into shape, steel items frequently go through a variety of heat treatments. Annealing is the process of heating them to 700–800 °C for a few hours, followed by a slow cooling. It softens and improves the workability of the steel.

Direct iron Reduction

There are no alternate ways to process iron due to environmental concerns. “Direct iron reduction” is the process of reducing iron ore into a ferrous lump that may be used to make steel, known as “sponge” iron or “direct” iron. The direct reduction process consists of two primary reactions:
Heat and a catalyst are used to partially oxidize natural gas:

{2 CH₄ + O₂ -> 2 CO + 4H₂}
After that, these gasses are applied to iron ore in a furnace to create solid sponge iron:
{Fe₂O₃ + CO + 2H₂ -> 2Fe + CO₂ + 2H₂O}
As previously mentioned, silica is eliminated by adding a limestone flux.

Thermite Process

Through the thermite reaction, metallic iron is produced when a mixture of iron oxide and aluminum powder ignites:
{Fe₂O₃ + 2 Al -> 2 Fe + Al₂O₃}
Alternatively, pig iron can be processed to create wrought iron, which is commercially pure iron, or steel, which contains up to 2% carbon.

This has been accomplished by a variety of techniques, such as electric arc furnaces, puddling furnaces, Bessemer converters, open hearth furnaces, basic oxygen furnaces, and finery forges.

The goal is always to oxidize part or all of the carbon together with other contaminants. On the other hand, alloy steels can be created by adding additional metals.

Molten Oxide Electrolysis

An alloy of iron, chromium, and other metals that does not react with oxygen is used in molten oxide electrolysis, together with a liquid iron cathode.

The electrolyte is a combination of molten metal oxides into which iron ore is dissolved. While lowering the iron oxide, the current maintains the electrolyte’s molten state.

Apart from the pure liquid iron, oxygen is also generated and can be sold to somewhat recoup the expenses.

The size of production cells varies and might be substantially lower than that of traditional furnaces. The electricity required to heat and reduce the metal is the only source of carbon dioxide emissions.

Applications

As Structural Material

With over 90% of all metal production occurring worldwide, iron is the most frequently utilized metal.

It is frequently chosen as the material of choice to bear stress or transfer pressures because of its inexpensive cost and high strength.

Examples of applications include the fabrication of machinery and machine tools, trains, cars, ship hulls, concrete reinforcing bars, and the load-carrying structure of buildings. Iron is usually mixed with alloying substances to create steel, as pure iron is rather soft.

Material	TS (MPa)	BH (Brinell)
Iron whiskers	11000
Ausformed (hardened) steel	2930	850–1200
Martensitic steel	2070	600
Bainitic steel	1380	400
Pearlitic steel	1200	350
Cold-worked iron	690	200
Small-grain iron	340	100
Carbon-containing iron	140	40
Pure, single-crystal iron	10	3

Mechanical Properties

The structural applications of iron and its alloys are highly dependent on their mechanical properties.

Numerous tests, such as the Brinell, Rockwell, and Vickers hardness tests, can be used to assess those qualities.

It is common practice to compare tests or calibrate measurements using the characteristics of pure iron.

The purity of the sample, however, has a major impact on the mechanical characteristics of iron: pure, single crystals of iron are actually softer than aluminum, and the finest iron produced industrially (99.99%) has a hardness of 20–30 Brinell.

Electrolytic refining is an industrial process used to create pure iron (99.9%?99.999%), also known as electrolytic iron.

Iron’s hardness and tensile strength will rise significantly with an increase in carbon content. Despite the alloy’s low tensile strength, a 0.6% carbon content allows for a maximum hardness of 65 Rc.

Iron is significantly easier to work with than its heavier counterparts, ruthenium and osmium, due to its softness.

Types of Steels and Alloys

At 910°C, a-Iron, a rather soft metal, may dissolve a small amount of carbon (maximum 0.021% by mass).

Similar to iron, austenite may dissolve a significant amount of carbon (up to 2.04% by mass at 1146 °C), yet it is still soft and metallic.

This sort of iron is utilized to make stainless steel, which is used to make cutlery and equipment for hospitals and food services.

The iron that is sold commercially is categorized according to its level of purity and quantity of additions.

Pig iron has a carbon content of 3.5–4.5% with varied concentrations of impurities like silicon, phosphorus, and sulfur.

Pig iron is a byproduct of the production of cast iron and steel; it is not a product that can be sold.

The removal of impurities from pig iron that have an adverse effect on the material’s qualities, produces cast iron, which has 2-4 percent carbon, 1-6 percent silicon, and trace amounts of manganese.

As sulfur and phosphorus, When iron and carbon are heated together, pig iron melts first because its melting point is between 1420 and 1470K, which is lower than that of either of its two constituents.

The shape that the carbon acquires in the alloy determines the wide range of mechanical properties that it has.

Iron carbide (Fe₃C), often known as cementite, is the form of carbon found in “white” cast irons. White cast irons’ mechanical characteristics are dominated by this brittle, hard substance, which makes them strong but susceptible to shock.

The name comes from the fine facets of fractured iron carbide, a very pale, silvery, glossy substance that is present on the broken surface of white cast irons.

lowering a combination of iron and 0.8% carbon gradually to room temperature below 723 °C The temperature causes distinct, alternating layers of a-iron, often known as pearlite due to its appearance, and cementite, which is soft and flexible.

On the other hand, rapid cooling produces hard and brittle martensite by preventing time for this separation.

After that, the steel can be tempered by heating it to a temperature that is intermediate, which will alter the pearlite and martensite ratios.

The final product has a pearlite-aFe mixture below 0.8% carbon content and a pearlite-cementite mixture above 0.8% carbon content.

The carbon in gray iron is present as discrete, fine graphite flakes. The brittle nature of the material is caused by the sharp edges of the graphite flakes, which create locations where stress concentrates.

A more recent variety of gray iron, called ductile iron, is carefully treated with tiny amounts of magnesium to change the graphite’s shape to spheroids, or nodules.

This reduces stress concentrations and significantly improves the material’s toughness and strength.

Less than 0.25% of wrought iron is carbon, yet a significant proportion of slag gives it a fibrous texture.

Though not as fusible as pig iron, it is a sturdy and pliable product. It quickly loses its edge if it is sharpened.

One distinguishing feature of wrought iron is the tiny slag threads that are trapped inside the metal. Compared to steel, wrought iron is more resistant to corrosion.

For conventional “wrought iron” items, mild steel has nearly entirely superseded it, as has blacksmithing.

Although mild steel is more affordable and accessible than wrought iron, it corrodes more easily. Less than 2.0% of carbon is present in carbon steel, together with trace amounts of silicon, sulfur, phosphorus, and manganese.

Different proportions of carbon are included in alloy steels, along with additional metals including chromium, vanadium, molybdenum, nickel, tungsten, etc.

Because of their increased cost due to their alloy content, they are typically exclusively used in specialized applications.

However, stainless steel is a typical alloy steel. An expanding variety of micro-alloyed steels, often known as “HSLA” or high-strength, low-alloy steels, have been created thanks to advancements in ferrous metallurgy recently.

These steels are made with minute additions that result in high strengths and frequently amazing toughness at low cost.

In particular, alloys with high-purity elemental compositions—such as electrolytic iron alloys—have improved qualities like toughness, fatigue strength, ductility, tensile strength, heat resistance, and corrosion resistance.

In addition to its conventional uses, iron also serves as an ionizing radiation shield. It is mechanically far stronger than lead, another conventional protective substance, although being lighter. The graph displays radiation attenuation as a function of energy.

The primary drawback of iron and steel is that, if left unprotected, pure iron and the majority of its alloys rust badly, costing more than 1% of global GDP.

Iron can be prevented from rusting by cathodic protection, painting, galvanization, passivation, plastic coating, and bluing, which all work by keeping out oxygen and water. Here is the mechanism by which iron rusts:

Cathode: 3 O₂ + 6 H₂O + 12 e⁻ → 12 OH⁻

Anode: 4 Fe → 4 Fe²⁺ + 8 e⁻; 4 Fe²⁺ → 4 Fe³⁺ + 4 e⁻

Overall: 4 Fe + 3 O₂ + 6 H₂O → 4 Fe³⁺ + 12 OH⁻ → 4 Fe(OH)₃ or 4 FeO(OH) + 4 H₂O
In metropolitan locations, iron(II) sulfate is often generated when air sulfur dioxide attacks iron, while in seaside areas, atmospheric salt particles serve as the electrolyte.

Catalysts and Reagents

Fe is cheap and harmless, which is why a lot of work has gone into creating catalysts and reagents based on Fe.

However, compared to more costly metals, iron is used less frequently as a catalyst in industrial operations. Fe-containing enzymes are widely found in biology.

The Fischer-Tropsch process, which turns carbon monoxide into hydrocarbons used in fuels and lubricants, and the Haber-Bosch process, which produces ammonia, have historically both included iron catalysts.

Nitrobenzene is converted to aniline by a process called Bechamp reduction, which uses powdered iron in an acidic medium.

Iron Compounds

Thermite reactions can be produced by lighting iron(III) oxide and aluminum powder mixtures, which are then utilized to purify ore and weld big iron components like rails.

Oxyhydroxide and iron(III) oxide are utilized as ocher and reddish pigments. Iron(III) chloride is used as an etchant for copper in printed circuit board manufacturing,

As an additive in animal feed, as a coloring agent in paints, for fabric dyeing, water purification, and sewage treatment.

Iron tincture, a medication used to reduce canary bleeding, can also be made by dissolving it in alcohol.

One uses iron(II) sulfate as a starting point for additional iron compounds. It’s also applied to cement to lessen chromate.

It is used to treat iron-deficient anemia and enrich meals. Tank water is treated with iron(III) sulfate to settle minute sewage particles.

In the production of iron complexes and magnetic iron oxides, as well as in organic synthesis, iron(II) chloride is employed as a reducing flocculating agent.

One medication used as a vasodilator is sodium nitroprusside. It is included in the List of Essential Medicines by the World Health Organization.

Biological and Pathological Role

Life needs iron. The enzymes known as nitrogenase, which are in charge of biological nitrogen fixation, are part of the ubiquitous iron-sulfur clusters.

Proteins containing iron are involved in the usage, storage, and transportation of oxygen. Electron transport involves iron proteins.

Higher species have proteins that include iron, such as catalase, cytochrome (see high-valent iron), and hemoglobin.

Because the human body recycles its hemoglobin for iron content, the average adult human contains about 0.005% body weight of iron, or about four grams, of which three-quarters is in hemoglobin.

This level of iron remains constant despite only about one milligram of iron being absorbed each day. Iron (II) can be reduced or oxidized to promote microbial growth (III).

Biochemistry

For aerobic species, iron uptake presents a challenge since ferric iron is weakly soluble at pH neutral.

As a result, these organisms have evolved mechanisms for absorbing iron in the form of complexes. Occasionally, they will take up ferrous iron and oxidize it back to ferric iron.

Particularly, bacteria have developed siderophores, which are exceptionally high-affinity sequestering agents.

Iron accumulation in human cells is carefully controlled after uptake. The protein transferrin, which binds iron ions taken up from the duodenum and transports them to cells via the blood, is a key player in this control.

With a high stability constant, transferrin is highly efficient at absorbing Fe³⁺ ions from even the most stable complexes. Fe³⁺ is attached to one nitrogen, three oxygens, and a chelating carbonate anion that traps the Fe³⁺ ion in the midst of a deformed octahedron.

Transferrin is converted from Fe³⁺ and Fe²⁺ to ferritin in the bone marrow in order to be integrated into hemoglobin.

Among the most well-known and researched bioinorganic iron compounds, or biological iron molecules, are the heme proteins, which include cytochrome P450, myoglobin, and hemoglobin.

These substances take a role in the manufacturing of enzymes, the transportation of gasses, and the transfer of electrons.

A class of proteins known as “metalloproteins” has cofactors that are metal ions. Iron metalloproteins include, for instance, ferritin and rubredoxin.

Iron is found in many enzymes that are essential to life, including lipoxygenases, IRE-BP, and catalase.

The oxygen carrier hemoglobin, which gives red blood cells their color, carries oxygen from the lungs down the arteries to the muscles.

Myoglobin then stores the oxygen until it is required for the metabolic oxidation of glucose, which produces energy.

Here, hemoglobin attaches itself to carbon dioxide, which is created when glucose is oxidized and is carried by hemoglobin (mostly as bicarbonate anions) back to the lungs for exhalation.

Hemoglobin has six potential coordination sites, with the iron located in one of the four heme groups.

Four of these sites are filled by nitrogen atoms in porphyrin rings, while the fifth is by imidazole nitrogen in a histidine residue that is connected to one of the protein chains.

The oxygen molecule it can reversibly attach to is assigned the sixth position, which is reserved for the heme group.

The Fe²⁺ ion at the center of the heme group, in the interior of the hydrophobic protein, is in a high-spin state when hemoglobin is not connected to oxygen (and is referred to as deoxyhemoglobin).

Due to its size, it cannot fit within the porphyrin ring; instead, it bends into a dome, with the Fe²⁺ ion situated approximately 55 picometers above it.

In this arrangement, another block the sixth coordination site set out for the oxygen remnant of tidiness.

This histidine residue goes away from deoxyhemoglobin when it picks up an oxygen molecule and returns when the oxygen molecule is firmly bonded to establish a hydrogen bond with it.

The Fe²⁺ ion switches to a low-spin state as a result, causing the ionic radius to fall by 20% and making it fit inside the planar porphyrin ring.

Moreover, the oxygen molecule is tilted as a result of this hydrogen bonding, creating a Fe–O–O bond angle of about 120° that prevents the creation of Fe–O–Fe or Fe–O₂–Fe bridges.

Which might result in electron transfer, the oxidation of Fe²⁺ to Fe^3+, as well as the breakdown of hemoglobin.) All of the protein chains migrate as a result, causing the other hemoglobin subunits to change shape and take on a bigger oxygen affinity.

Deoxyhemoglobin’s propensity for additional oxygen so rises when it absorbs oxygen, and vice versa.

However, because myoglobin only has one heme group, this cooperative effect is not possible.

Therefore, even though hemoglobin is nearly saturated with Compared to myoglobin, which oxygenates even at low partial pressures of oxygen present in muscle tissue, its affinity for oxygen is significantly lower at the high partial pressures of oxygen found in the lungs.

The Bohr effect, named after Niels Bohr’s father Christian Bohr, states that the presence of carbon dioxide reduces hemoglobin’s affinity for oxygen.

Because they bind to hemoglobin similarly to oxygen but much more strongly, carbon monoxide and phosphorus trifluoride are toxic to humans because they prevent oxygen from being carried throughout the body.

Carboxyhemoglobin is hemoglobin that has been bonded to carbon monoxide. This effect also contributes slightly to cyanide’s toxicity, but its main impact is unquestionably its disruption of cytochrome a, the electron transport protein, from performing as intended.

Heme groups are also involved in the metabolic oxidation of glucose by oxygen, which is carried out by the cytochrome proteins.

Next, one of two entities occupies the sixth coordination site: Except for cytochrome a, which binds directly to oxygen and is thus rapidly contaminated by cyanide, these proteins are mostly inert to oxygen due to the presence of another imidazole nitrogen or methionine sulfur.

Here, the iron maintains its low spin while transitioning between the +2 and +3 oxidation states, allowing for the electron transfer to occur.

Because each step’s reduction potential is marginally higher than the one before it, energy is released gradually and can be stored in adenosine triphosphate.

A little different is cytochrome a, which is found in the mitochondrial membrane, attaches directly to oxygen, and moves protons in addition to electrons as follows:

4 Cytc²⁺ + O₂ + 8H⁺_inside → 4 Cytc³⁺ + 2 H₂O + 4H⁺_outside

While heme proteins are the most significant group of proteins containing iron, iron-sulfur proteins are also highly significant because they play a role in electron transfer.

This is made feasible by the fact that iron can reside steadily in both the +2 and +3 oxidation states.

These consist of one, two, four, or eight iron atoms that are roughly tetrahedrally coupled to four sulfur atoms; as a result, they are invariably high-spin iron.

The One iron atom coordinated with four sulfur atoms from cysteine residues in the surrounding peptide chains makes up rubredoxin, the most basic of these molecules.

Ferredoxins are a significant subclass of iron-sulfur proteins that contain several iron atoms. Transferrin is not a member of any of these groups mussels’ utilization of organometallic iron-based linkages in their protein-rich cuticles helps them to hold onto rocks in the water.

The iron content of these structures was found to improve toughness by 92 times, tensile strength by 58 times, and elastic modulus by 770 times, based on synthetic reproductions. The stress needed to harm them irreversibly rose by 76 times.

Nutrition

Diet

Although iron is found in many foods, red meat, oysters, beans, chicken, fish, leafy greens, watercress, tofu, and blackstrap molasses are especially rich dietary iron sources.

There are occasions when iron is added especially to bread and cereal for breakfast.Iron(II) fumarate is the most common form of iron found in dietary supplements; however, iron(II) sulfate is less expensive and similarly well absorbed.

Even though it is only absorbed at a rate of one-third to two-thirds that of iron sulfate, elemental iron, also known as reduced iron, is frequently added to foods like enriched wheat flour and morning cereals.

The body can absorb iron best when it is chelated to amino acids. It can also be taken as a regular iron supplement. The least expensive amino acid, glycine, is most frequently utilized in the production of iron glycinate supplements.

Dietary Recommendations

The 2001 revision of the Recommended Dietary Allowances (RDAs) and Estimated Average Requirements (EARs) for iron was made by the U.S. Institute of Medicine (IOM).

The current EAR for iron is 7.9 mg/day for women ages 14 to 18, 8.1 mg/day for women ages 19 to 50, and 5.0 mg/day for women ages 50 and beyond (post-menopause).

For males aged 19 and above, the EAR is 6.0 mg/day. 15.0 mg/day for women aged 15 to 18, 18.0 mg/day for those aged 19 to 50, and 8 upward.

RDAs are higher than EARs to identify levels that will cover people with needs that are above average.

RDAs are 27 mg/day for pregnant women and 9 mg/day for nursing mothers. 7 mg/day for kids 1-3 years old, 10 mg/day for ages 4-8, and 8 mg/day for ages 9-13.

When there is enough data to support a claim, the IOM also establishes acceptable upper intake levels (ULs) for vitamins and minerals in terms of safety.

The recommended daily allowance (UL) for iron is 45 mg. Dietary Reference Intakes is the term used to refer to the combined values of EARs, RDAs, and ULs.

The combined collection of data is referred to as Dietary Reference Values by the European Food Safety Authority (EFSA), which uses Population Reference Intake (PRI) rather than RDA and Average Requirement (EAR) in place of EAR.

The US definitions of AI and UL apply here as well. The recommended daily intake (PRI) for women is 13 mg for those aged 15 to 17, 16 mg for premenopausal women aged 18 and older, and 11 mg for postmenopausal women.

16 milligrams per day throughout nursing and pregnancy. The PRI for males aged 15 and up is 11 mg/day.

The PRI rises from 7 to 11 mg/day for kids aged 1 to 14. With the exception of pregnancy, the PRIs are greater than the RDAs in the United States.

The same safety question was investigated by the EFSA, but no UL was established. If babies are given cow’s milk in a bottle, they could need iron supplements.

Regular blood donors are frequently told to increase their iron intake since they run the risk of having low iron levels.

To comply with U.S. regulations for food and dietary supplement labeling, the quantity in a serving is stated as a percentage of Daily Value (%DV).

100% of the Daily Value for iron labeling was 18 mg, and as of May 27, 2016, that amount had not changed. There is a table with the previous and current adult daily values.

Deficiency

Worldwide, iron deficiency is the most common nutritional deficiency. A condition known as latent iron deficiency arises when iron loss is not sufficiently offset by adequate dietary iron intake.

If iron deficiency anemia is not treated, this condition eventually results in inadequate red blood cells and hemoglobin.

The condition is particularly common in children, women who are not yet menopausal, and those who have bad diets.

Although the majority of cases of iron-deficiency anemia are minor, if left untreated, it can lead to issues during pregnancy, fast or irregular heartbeats, and delayed growth in children and babies.

The brain’s blood-brain barrier prevents acute iron insufficiency because iron moves slowly through it.

Serum ferritin levels, which indicate acute variations in iron status, may not accurately reflect brain iron status; nonetheless, chronic dietary iron insufficiency may gradually lower brain iron concentrations.

Iron is involved in the transport of oxygen, the production of myelin, mitochondrial respiration, and the creation and metabolism of neurotransmitters in the brain.

Biomolecular alterations similar to those observed in Parkinson’s and Huntington’s disease have been reported in animal models of dietary iron deficiency.

On the other hand, Parkinson’s disease has also been connected to age-related iron buildup in the brain.

Excess

The human body has a strict regulation on iron intake because it lacks a regulated physiological mechanism for iron excretion.

Because mucosal and skin epithelial cell sloughing only results in minute losses of iron, the main method of controlling iron levels is uptake regulation.

A genetic mutation that corresponds to the HLA-H gene area on chromosome 6 causes unusually low amounts of hepcidin, a critical regulator of iron entrance into the circulatory system in animals, and impairs the regulation of iron uptake in some individuals.

In certain individuals, high iron consumption may lead to iron overload diseases, also referred to as hemochromatosis in medicine.

Many people are unaware of a family history of iron overload and have an undetected hereditary vulnerability to the condition.

Because of this, people shouldn’t use iron supplements unless they have been diagnosed with iron deficiency and have seen a physician.

It is estimated that between 0.3% and 0.8% of all metabolic illnesses in Caucasians are caused by hemochromatosis.

Excessive blood levels of free iron can result from iron overdosing. Elevated quantities of free ferrous iron in the blood combine with peroxides to form very reactive free radicals that can harm lipids, proteins, DNA, and other parts of the cell.

When the amount of iron in the cell is greater than the capacity of transferrin to bind the iron, free iron is present and can lead to iron toxicity.

Damage to the gastrointestinal tract’s cells may also make it difficult for them to control the absorption of iron, which could result in additional blood level rises.

Usually, iron harms cells in the liver, heart, and other areas. resulting in detrimental consequences such as death, long-term organ damage, shock, liver failure, metabolic acidosis, coagulopathy, and coma.

Iron poisoning occurs in humans when the blood contains more iron than 20mg per kilogram of body weight; 60mg per kilogram is the threshold for death.

One of the most frequent toxicological causes of death in children under six is excessive iron consumption, which frequently results from kids consuming huge amounts of adult-only ferrous sulfate tablets.

For individuals, the Tolerable Upper Intake Level (UL) is 45mg/day based on the Dietary Reference Intake (DRI). The UL is 40 mg/day for kids under the age of 14.

The complex medical treatment of iron toxicity may involve binding and eliminating excess iron from the body with a particular chelating drug called deferoxamine.

ADHD

Low levels of thalamic iron have been linked to the pathophysiology of ADHD, according to some research.

Iron supplementation has been observed by some researchers to be beneficial, particularly in the inattentive subtype of the condition.

Additionally, iron may be able to lower the risk of cardiovascular events while using ADHD medications, according to one study.

In the 2000s, some studies hypothesized a connection between ADHD and low blood iron levels. No such association was discovered in a 2012 investigation.

Cancer

Iron plays a “double-edged sword” role in cancer defense since it is present in many non-pathological processes.

Intravenous iron therapy is used to restore iron levels in patients undergoing chemotherapy who develop anemia and iron shortage.

Iron excess, which can result from consuming a lot of red meat, can cause tumors to grow and make people more vulnerable to developing cancer, especially colon cancer.

Marine Systems

In marine environments, iron is vital and can function as a nutrient that limits planktonic activity. For this reason, a reduction in iron levels that is too great could cause a fall in the growth rates of diatoms and other phytoplanktonic species.

In environments with high iron and low oxygen, marine bacteria can also oxidize iron. Both directly from the atmosphere and through nearby rivers, iron can infiltrate marine systems.

Iron can be recycled at the cellular level and through ocean mixing spread throughout the water column once it reaches the ocean.

Sea ice in the Arctic is essential to the distribution and storage of iron in the ocean because it replenishes marine iron throughout the winter months when it freezes and returns it to the water during the summer.

The availability of iron to primary producers can change due to the iron cycle, which can change the shape of iron from aqueous to particulate forms.

There is more iron in forms that primary producers can use when there is more sunshine and warmth.

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