26
  Fe  
55.845000
Iron

Name: Iron
Symbol: Fe
Atomic Number: 26
AtomicWeight: 55.845000
Family: Transition metals
CAS RN: 7439-89-6
Description: Pure iron is lustrous, silvery and easy to work. Iron easily rusts in damp air.
State (25C): Solid
Oxidation states: +2, +3

Molar Volume: 7.11 cm3/mole
Valence Electrons: 3d64s2

Boiling Point:  3023K, 2750C, 4982F
Melting Point:
1808K, 1535C, 2795F
Electrons Energy Level: 2, 8, 14, 2
Isotopes: 24 + 4 Stable
Heat of Vaporization: 349.6 kJ/mol
Heat of Fusion: 13.8 kJ/mol
Density: 7.874 g/cm3 @ 300K
Specific Heat: 0.44 J/gK
Atomic Radius: 1.72
Ionic Radius: 0.645
Electronegativity: 1.83 (Pauling); 1.64 (Allrod Rochow)
Vapor Pressure: 7.05 Pa @ 1535C

1s2 2s2p6 3s2p6d6 4s2

History

The first iron used by mankind, far back in prehistory, came from meteors.  The smelting of iron in bloomeries, probably began in Anatolia or the Caucasus in the second millennium BC or the latter part of the preceding one.  Cast iron was first produced in China about 550 BC, but not in Europe until the medieval period.  During the medieval period, means were found in Europe of producing wrought iron from cast iron (in this context known as pig iron) using finery forges.  For all these processes, charcoal was required as fuel.

iron.gif (826 bytes)

Alchemical Symbol, Ferrum

Steel (with a smaller carbon content than pig iron but more than wrought iron) was first produced in antiquity.  New methods of producing it by carburizing bars of iron in the cementation process were devised in the 17th century AD.  In the Industrial Revolution,  new methods of producing bar iron without charcoal were devised and these were later applied to produce steel. In the late 1850s, Henry Bessemer invented a new steelmaking process, involving blowing air through molten pig iron, to produce mild steel.  This and other 19th century and later processes have led to wrought iron no longer being produced.

iron1.jpg (1147 bytes) iron2.jpg (1069 bytes) iron3.jpg (1158 bytes) iron4.jpg (1152 bytes)

Additional Representations of Alchemical Symbols for Iron

Characteristics

Iron is believed to be the tenth most abundant element in the universe, and fourth most abundant on earth.  The concentration of iron in the various layers in the structure of the Earth ranges from high (probably greater than 80%, perhaps even a nearly pure iron crystal) at the inner core, to only 5% in the outer crust.  Iron is second in abundance to aluminum among the metals and fourth in abundance in the crust.  Iron is the most abundant element by mass of our entire planet, making up 35% of the mass of the Earth as a whole.

1s2
2s2 2p6
3s2 3p6 3d6
4s2

Iron is a metal extracted from iron ore, and is almost never found in the free elemental state.  In order to obtain elemental iron, the impurities must be removed by chemical reduction.  Iron is the main component of steel, and it is used in the production of alloys or solid solutions of various metals, as well as some non-metals, particularly carbon.  The many iron-carbon alloys, which have very different properties, are discussed in the article on steel.

Nuclei of iron have some of the highest binding energies per nucleon, surpassed only by the nickel isotope 62Ni.  The universally most abundant of the highly stable nuclides is, however, 56Fe.  This is formed by nuclear fusion in stars.  Although a further tiny energy gain could be extracted by synthesizing 62Ni, conditions in stars are unsuitable for this process to be favoured, and iron abundance on Earth greatly favors iron over nickel, and also presumably in supernova element production.  When a very large star contracts at the end of its life, internal pressure and temperature rise, allowing the star to produce progressively heavier elements, despite these being less stable than the elements around mass number 60, known as the "iron group".  This leads to a supernova.

Some cosmological models with an open universe predict that there will be a phase where as a result of slow fusion and fission reactions, everything will become iron.

Iron (as Fe2+, ferrous ion) is a necessary trace element used by all known living organisms.  Iron-containing enzymes, usually containing heme prosthetic groups, participate in catalysis of oxidation reactions in biology, and in transport of a number of soluble gases.

Occurrence

180px-IronInRocksMakeRiverRed.jpg (10638 bytes)

The red appearance of this water is due to iron in the rocks.

Iron is one of the most common elements on Earth, making up about 5% of the Earth's crust.  Most of this iron is found in various iron oxides, such as the minerals hematite, magnetite, and taconite.  The earth's core is believed to consist largely of a metallic iron-nickel alloy.  About 5% of the meteorites similarly consist of iron-nickel alloy. Although rare, these are the major form of natural metallic iron on the earth's surface.

The reason for Mars's red color is thought to be an iron-rich soil.

Iron Ore Production

Industrially, iron is produced starting from iron ores, principally hematite (nominally Fe2O3) and magnetite (Fe3O4) by a carbothermic reaction (reduction with carbon) in a blast furnace at temperatures of about 2000C.   In a blast furnace, iron ore, carbon in the form of coke, and a flux such as limestone are fed into the top of the furnace, while a blast of heated air is forced into the furnace at the bottom.

In the furnace, the coke reacts with oxygen in the air blast to produce carbon monoxide: 

6C + 3O2  rarrow.gif (63 bytes) 6CO

The carbon monoxide reduces the iron ore (in the chemical equation below, hematite) to molten iron, becoming carbon dioxide in the process:

6CO  + 2Fe2O3 rarrow.gif (63 bytes) 4Fe + 6CO2

The flux is present to melt impurities in the ore, principally silicon dioxide sand and other silicates.  Common fluxes include limestone (principally calcium carbonate) and dolomite (calcium-magnesium carbonate).  Other fluxes may be used depending on the impurities that need to be removed from the ore. In the heat of the furnace the limestone flux decomposes to calcium oxide (quicklime):

CaCO3  rarrow.gif (63 bytes) CaO  + CO2

Then calcium oxide combines with silicon dioxide to form a slag.

CaO + SiO2  rarrow.gif (63 bytes) CaSiO3

The slag melts in the heat of the furnace, which silicon dioxide would not have.   In the bottom of the furnace, the molten slag floats on top of the more dense molten iron, and apertures in the side of the furnace are opened to run off the iron and the slag separately.  The iron once cooled, is called pig iron, while the slag can be used as a material in road construction or to improve mineral-poor soils for agriculture.

Pig iron is not pure iron, but has 4-5% carbon dissolved in it.  This is subsequently reduced to steel or commercially pure iron, known as wrought iron, using other furnaces or converters.

Approximately 1100Mt (million tons) of iron ore was produced in the world in 2000, with a gross market value of approximately 25 billion dollars.  While ore production occurs in 48 countries, the five largest producers were China, Brazil, Australia, Russia and India, accounting for 70% of world iron ore production.  The 1100Mt of iron ore was used to produce approximately 572Mt of pig iron.

Applications

Iron is the most used of all the metals, comprising 95% of all the metal tonnage produced worldwide.  Its combination of low cost and high strength make it indispensable, especially in applications like automobiles, the hulls of large ships, and structural components for building.  Steel is the best known alloy of iron, and some of the forms that iron can take include:

The main drawback to iron and steel is that pure iron, and most of its alloys, suffer badly from rust if not protected in some way.  Painting, galvanization, plastic coating and bluing are some techniques used to protect iron from rust by excluding water and oxygen or by sacrificial protection.

Compounds

Isotopes

Naturally occurring iron consists of four isotopes: 5.845% of radioactive 54Fe (half-life: >3.11022 years), 91.754% of stable 56Fe, 2.119% of stable 57Fe and 0.282% of stable 58Fe.  60Fe is an extinct radionuclide of long half-life (1.5 million years).

Much of the past work on measuring the isotopic composition of Fe has centered on determining 60Fe variations due to processes accompanying nucleosynthesis (i.e., meteorite studies) and ore formation.  In the last decade however, advances in mass spectrometry technology have allowed the detection and quantification of minute, naturally-occurring variations in the ratios of the stable isotopes of iron.  Much of this work has been driven by the Earth and planetary science communities, although applications to biological and industrial systems are beginning to emerge.

The isotope 56Fe is of particular interest to nuclear scientists.  A common misconception is that this isotope represents the most stable nucleus possible, and that it thus would be impossible to perform fission or fusion on 56Fe and still liberate energy.  This is not true, as both 62Ni and 58Fe are more stable.

In phases of the meteorites Semarkona and Chervony Kut a correlation between the concentration of 60Ni, the daughter product of 60Fe, and the abundance of the stable iron isotopes could be found which is evidence for the existence of 60Fe at the time of formation of the solar system.  Possibly the energy released by the decay of 60Fe contributed, together with the energy released by decay of the radionuclide 26Al, to the remelting and differentiation of asteroids after their formation 4.6 billion years ago.  The abundance of 60Ni present in extraterrestrial material may also provide further insight into the origin of the solar system and its early history.  Of the stable isotopes, only 57Fe has a nuclear spin (-1/2).

atom.gif (700 bytes)

Isotope  Atomic Mass Half-Life
45Fe 45.01458 4.9 ms
46Fe 46.00081 9 ms
47Fe 46.99289 21.8 ms
48Fe 47.98050 44 ms
49Fe 48.97361 70 ms
50Fe 49.96299 155 ms
51Fe 50.956820 305 ms
52Fe 51.948114 8.275 hours
53Fe 52.9453079 8.51 minutes
54Fe 53.9396105 Stable
55Fe 54.9382934 2.737 years
56Fe 55.9349375 Stable
57Fe 56.9353940 Stable
58Fe 57.9332756 Stable
59Fe 58.9348755 44.495 days
60Fe 59.934072 1.5 x 106 years
61Fe 60.936745 5.98 minutes
62Fe 61.936767 68 seconds
63Fe 62.94037 6.1 seconds
64Fe 63.9412 2.0 seconds
65Fe 64.94538 1.3 seconds
66Fe 65.94678 440 ms
67Fe 66.95095 394 ms
68Fe 67.95370 187 ms
69Fe 68.95878 109 ms
70Fe 69.96146 94 ms
71Fe 70.96672 >300 ns
72Fe 71.96962 >300 ns

Biological Role

Organic Synthesis

The usage of iron metal filings in organic synthesis is mainly for the reduction of nitro compounds.  Additionally, iron has been used for desulfurizations, reduction of aldehydes, and the deoxygenation of amine oxides.

180px-Heme.svg.jpg (8508 bytes)

Structure of Heme b

Iron is essential to nearly all known organisms.  It is mostly stably incorporated in the inside of metalloproteins, because in exposed or in free form it causes production of free radicals that are generally toxic to  cells. To say that iron is free doesn't mean that it is free floating in the bodily fluids.  Iron binds avidly to virtually all biomolecules so it will adhere nonspecifically to cell membranes, nucleic acids, proteins, etc. 

Many animals incorporate iron into the heme complex, an essential component of cytochromes, which are proteins involved in redox reactions (including but not limited to cellular respiration), and of oxygen carrying proteins hemoglobin and myoglobin.   Inorganic iron involved in redox reactions is also found in the iron-sulfur clusters of many enzymes, such as nitrogenase (involved in the synthesis of ammonia from nitrogen and hydrogen) and hydrogenase.  A class of non-heme iron proteins is responsible for a wide range of functions within several life forms, such as enzymes methane monooxygenase (oxidizes methane to methanol), ribonucleotide reductase ( reduces ribose to deoxyribose; DNA biosynthesis), hemerythrins (oxygen transport and fixation in marine invertebrates) and purple acid phosphatase (hydrolysis of phosphate esters).    When the body is fighting a bacterial infection, the body sequesters iron inside of cells (mostly stored in the storage molecule ferritin) so that it cannot be used by bacteria.

Iron distribution is heavily regulated in mammals, both as a defense against bacterial infection and because of the potential biological toxicity of iron.  The iron absorbed from the duodenum binds to transferrin, and is carried by blood to different cells.  There There it gets incorporated, by an as yet unknown mechanism, into target proteins.

Nutrition & Dietary Sources

Good sources of dietary iron include meat, fish, poultry, lentils, beans, leaf vegetables, tofu, chickpeas, black-eyed peas, potatoes with skin, bread made from completely whole-grain flour, and molasses.

Iron provided by dietary supplement is often found as Iron (II) fumarate.  Iron sulfate is as well absorbed, and less expensive. Elemental iron, despite being absorbed to a much smaller extent, is often added to foods like breakfast cereals or "enriched" wheat flour (and will be listed as "reduced iron" in the list of ingredients).  The most bioavailable form of iron supplement (ten to fifteen times more bioavailable than any other) is iron amino acid chelate.  The RDA for iron varies considerably based on the age, gender, and source of dietary iron (heme-based iron has higher bioavailability.

Precautions

40px-Skull_and_crossbones.svg.jpg (1420 bytes) Excessive iron is toxic to humans, because excess ferrous iron reacts with peroxides in the body, producing free radicals.  Iron becomes toxic when it exceeds the amount of transferrin needed to bind free iron.  In excess, uncontrollable quantities of free radicals are produced.

Iron uptake is tightly regulated by the human body, which has no physiologic means of excreting iron and regulates iron solely by regulating uptake.  However, too much ingested iron can damage the cells of the gastrointestinal tract directly, and may enter the bloodstream by damaging the cells that would otherwise regulate its entry.  Once there, it causes damage to cells in the heart, liver and elsewhere.  This can cause serious problems, including the potential of death from overdose, and long-term organ damage in survivors.

Humans experience iron toxicity above 20 milligrams of iron for every kilogram of weight, and 60 milligrams per kilogram is a lethal dose.  Over-consumption of iron, often the result of children eating large quantitities of ferrous sulfate tablets intended for adult consumption, is the most common toxicological cause of death in children under six.  The DRI lists the Tolerable Upper Intake Level (UL) for adults as 45 mg/day.   For children under fourteen years old the UL is 40 mg/day.

If iron intake is excessive in the context of a genetic predisposition iron overload disorders can sometimes result, such as hemochromatosis.  This has been mapped to the HLA-H gene region on chromosome 6.  Iron overload disorders require a genetic inability to regulate iron uptake; however, many people have a genetic susceptibility to iron overload without realizing it and without knowing a family history of the problem.   For this reason, people should not take iron supplements unless they suffer from iron deficiency and have consulted a doctor.  Blood donors are at special risk of low iron levels and are often advised to supplement their iron intake.  Hemochromatosis is estimated to cause disease in 0.3-0.8 percent of white people.

The medical management of iron toxicity is complex.  One element of the medical approach is a specific chelating agent called deferoxamine, used to bind and expel excess iron from the body in case of iron toxicity.

atom.gif (700 bytes)

Iron Data
 

Atomic Structure

  • Atomic Radius: 1.72
  • Atomic Volume: 7.1cm3/mol
  • Covalent Radius: 1.17
  • Cross Section (Thermal Neutron Capture) Barns: 2.56
  • Crystal Structure: Cubic body centered
  • Electron Configuration:
    1s2 2s2p6 3s2p6d6 4s2
  • Electrons per Energy Level: 2, 8, 14, 2
  • Ionic Radius: 0.645
  • Filling Orbital: 3d6
  • Number of Electrons (with no charge): 26
  • Number of Neutrons (most common/stable nuclide): 30
  • Number of Protons: 26
  • Oxidation States: 2, 3
  • Valence Electrons: 3d6 4s2

Chemical Properties

  • Electrochemical Equivalent: 0.69455 g/amp-hr
  • Electron Work Function: 4.7eV
  • Electronegativity: 1.83 (Pauling); 1.64 (Allrod Rochow)
  • Heat of Fusion: 13.8 kJ/mol
  • Incompatibilities:
  • Ionization Potential
    • First: 7.87
    • Second: 16.18
    • Third: 30.651
  • Valence Electron Potential (-eV): 67

Physical Properties

  • Atomic Mass Average: 55.847
  • Boiling Point: 3023K, 2750C, 4982F
  • Coefficient of Lineal Thermal Expansion/K-1: 12.3E-6
  • Conductivity
    Electrical: 0.0993 106/cm
    Thermal: 0.802 W/cmK
  • Density: 7.874 g/cm3 @ 300K
  • Description:
    Pure iron is lustrous, silvery and easy to work. Iron easily rusts in damp air.
  • Elastic Modulus:
    • Bulk: 170/GPa
    • Rigidity: 82/GPa
    • Youngs: 211/GPa
  • Enthalpy of Atomization: 414.2 kJ/mole @ 25C
  • Enthalpy of Fusion: 14.9 kJ/mole
  • Enthalpy of Vaporization: 351 kJ/mole
  • Flammablity Class:
  • Freezing Point: see melting point
  • Hardness Scale
    • Brinell: 490 MN m-2
    • Mohs: 4
    • Vickers: 608 MN m-2
  • Heat of Vaporization: 349.6 kJ/mol
  • Melting Point: 1808K, 1535C, 2795F
  • Molar Volume: 7.11 cm3/mole
  • Optical Reflectivity: 65%
  • Physical State (at 20C & 1atm): Solid
  • Specific Heat: 0.44 J/gK
  • Vapor Pressure: 7.05 Pa @ 1535C

Regulatory / Health

  • CAS Number
    • 7439-89-6
  • OSHA Permissible Exposure Limit (PEL)
    • No limits set by OSHA
  • OSHA PEL Vacated 1989
    • No limits set by OSHA
  • NIOSH Recommended Exposure Limit (REL)
    • No limits set by NIOSH
  • Levels In Humans:
    Note: this data represents naturally occuring levels of elements in the typical human, it DOES NOT represent recommended daily allowances.
    • Blood/mg dm-3: 447
    • Bone/p.p.m: 3-380
    • Liver/p.p.m: 250-1400
    • Muscle/p.p.m: 180
    • Daily Dietary Intake: 6-40 mg
    • Total Mass In Avg. 70kg human: 4.2 g
  • Discovery Year: Unknown
  • Name Origin:
    Latin, ferrum; Anglo-Saxon, iron
  • Abundance:
    • Earth's Crust/p.p.m.: 41000
    • Seawater/p.p.m.:
      • Atlantic Suface: 0.0001
      • Atlantic Deep: 0.0004
      • Pacific Surface: 0.00001
      • Pacific Deep: 0.0001
    • Atmosphere/p.p.m.: N/A
    • Sun (Relative to H=1E12): 3.16E+07
  • Sources:
    Obtained from hematite, magnetite, goethite, lepidocrocite and siderite. Annual world production is around 716,000,000 tons. Primary areas iron is mined are USA, Canada, Sweden, South Africa, Russia, India and Japan.
  • Uses:
    Used in steel and other alloys which are used in countless products. It is essential for animals as it is the chief constituent of hemoglobin which carries oxygen in blood vessels. Iron is the most important element of all the metals.
  • Additional Notes:
    Deficiency of iron leads to anaemia, but excess iron in the body causes liver and kidney damage.

Ionization Energy: 7.902 eV
Estimated Crustal Abundance: 5.63104 milligrams per kilogram
Estimated Oceanic Abundance: 210-3 milligrams per liter

Transition Metals
Group 3
(IIIB)
4
(IVB)
5
(VB)
6
(VIB)
7
(VIIB)
8
(VIIIB)
9
(VIIIB)
10 (VIIIB) 11
(IB)
12
(IIB)
Period 4 21
Sc
44.95
22
Ti
47.86
23
V
50.94
24
Cr
51.99
25
Mn
54.93
26
Fe
55.84
27
Co
58.93
28
Ni
58.69
29
Cu
63.54
30
Zn
65.39
Period 5 39
Y
88.90
40
Zr
91.22
41
Nb
92.90
42
Mo
95.94
43
Tc
98.00
44
Ru
101.0
45
Rh
102.9
46
Pd
106.4
47
Ag
107.8
48
Cd
112.4
Period 6 57
La
138.9
72
Hf
178.4
73
Ta
180.9
74
W
183.8
75
Re
186.2
76
Os
190.2
77
Ir
192.2
78
Pt
195.0
79
Au
196.9
80
Hg
200.5
Period 7 89
Ac
227.0
104
Rf
261.0
105
Db
262.0
106
Sg
266.0
107
Bh
264.0
108
Hs
269.0
109
Mt
268.0
110
Ds
269.0
111
Rg
272.0
112
Uub
277.0