| Discovery date | approx 3500BC |
| Discovered by | - |
| Origin of the name | The name comes from the Anglo-Saxon name 'iren'. |
| Allotropes |
|
Fe
Iron
26
55.845
| Group | 8 | Melting point | 1538°C, 2800°F, 1811 K |
| Period | 4 | Boiling point | 2861°C, 5182°F, 3134 K |
| Block | d | Density (g cm−3) | 7.87 |
| Atomic number | 26 | Relative atomic mass | 55.845 |
| State at 20°C | Solid | Key isotopes | 56Fe |
| Electron configuration | [Ar] 3d64s2 | CAS number | 7439-89-6 |
| ChemSpider ID | 22368 | ChemSpider is a free chemical structure database | |
Image explanation
The image is of the alchemical symbol for iron. The symbol is shown against a rusty mild steel plate.
Appearance
A shiny, greyish metal that rusts in damp air.
Uses
Iron is an enigma – it rusts easily, yet it is the most important of all metals. 90% of all metal that is refined today is iron.
Most is used to manufacture steel, used in civil engineering (reinforced concrete, girders etc) and in manufacturing.
There are many different types of steel with different properties and uses. Ordinary carbon steel is an alloy of iron with carbon (from 0.1% for mild steel up to 2% for high carbon steels), with small amounts of other elements.
Alloy steels are carbon steels with other additives such as nickel, chromium, vanadium, tungsten and manganese. These are stronger and tougher than carbon steels and have a huge variety of applications including bridges, electricity pylons, bicycle chains, cutting tools and rifle barrels.
Stainless steel is very resistant to corrosion. It contains at least 10.5% chromium. Other metals such as nickel, molybdenum, titanium and copper are added to enhance its strength and workability. It is used in architecture, bearings, cutlery, surgical instruments and jewellery.
Cast iron contains 3–5% carbon. It is used for pipes, valves and pumps. It is not as tough as steel but it is cheaper. Magnets can be made of iron and its alloys and compounds.
Iron catalysts are used in the Haber process for producing ammonia, and in the Fischer–Tropsch process for converting syngas (hydrogen and carbon monoxide) into liquid fuels.
Biological role
Iron is an essential element for all forms of life and is non-toxic. The average human contains about 4 grams of iron. A lot of this is in haemoglobin, in the blood. Haemoglobin carries oxygen from our lungs to the cells, where it is needed for tissue respiration.
Humans need 10–18 milligrams of iron each day. A lack of iron will cause anaemia to develop. Foods such as liver, kidney, molasses, brewer’s yeast, cocoa and liquorice contain a lot of iron.
Natural abundance
Iron is the fourth most abundant element, by mass, in the Earth’s crust. The core of the Earth is thought to be largely composed of iron with nickel and sulfur.
The most common iron-containing ore is haematite, but iron is found widely distributed in other minerals such as magnetite and taconite.
Commercially, iron is produced in a blast furnace by heating haematite or magnetite with coke (carbon) and limestone (calcium carbonate). This forms pig iron, which contains about 3% carbon and other impurities, but is used to make steel. Around 1.3 billion tonnes of crude steel are produced worldwide each year.
Iron objects have been found in Egypt dating from around 3500 BC. They contain about 7.5% nickel, which indicates that they were of meteoric origin.
The ancient Hittites of Asia Minor, today’s Turkey, were the first to smelt iron from its ores around 1500 BC and this new, stronger, metal gave them economic and political power. The Iron Age had begun. Some kinds of iron were clearly superior to others depending on its carbon content, although this was not appreciated. Some iron ore contained vanadium producing so-called Damascene steel, ideal for swords.
The first person to explain the various types of iron was René Antoine Ferchault de Réaumur who wrote a book on the subject in 1722. This explained how steel, wrought iron, and cast iron, were to be distinguished by the amount of charcoal (carbon) they contained. The Industrial Revolution which began that same century relied extensively on this metal.
| Atomic radius, non-bonded (Å) | 2.04 | Covalent radius (Å) | 1.24 |
| Electron affinity (kJ mol−1) | 14.569 |
Electronegativity (Pauling scale) |
1.83 |
|
Ionisation energies (kJ mol−1) |
1st
762.466
2nd
1561.876
3rd
2957.469
4th
5287.4
5th
7236
6th
9561.7
7th
12058.74
8th
14575.08
|
||
| Common oxidation states | 6, 3, 2, 0, -2 | ||||
| Isotopes | Isotope | Atomic mass | Natural abundance (%) | Half life | Mode of decay |
| 54Fe | 53.940 | 5.845 | > 3.1 x 1022 y | EC-EC | |
| 56Fe | 55.935 | 91.754 | - | - | |
| 57Fe | 56.935 | 2.119 | - | - | |
| 58Fe | 57.933 | 0.282 | - | - | |
|
|
|
Specific heat capacity (J kg−1 K−1) |
449 | Young's modulus (GPa) | 211.4 (soft); 152.3 (cast) | |||||||||||
| Shear modulus (GPa) | 81.6 (soft); 60.0 (cast) | Bulk modulus (GPa) | 169.8 | |||||||||||
| Vapour pressure | ||||||||||||||
| Temperature (K) |
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| Pressure (Pa) |
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| Listen to Iron Podcast |
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Transcript :
Chemistry in its element: iron (Promo) You're listening to Chemistry in its element brought to you by Chemistry World, the magazine of the Royal Society of Chemistry. (End promo) Chris Smith Hello, this week we turn to one of the most important elements in the human body. It's the one that makes metabolism possible and don't we just know it. There are iron man challenges, iron fisted leaders and those said to have iron in the soul. But there's a dark side to element number 26 too because its powerful chemistry means that it's also bad news for brain cells as Nobel Laureate Kary Mullis explains Kary Mullis For the human brain, iron is essential yet deadly. It exists on Earth mainly in two oxidation states - FeII and FeIII. FeIII is predominant within a few meters of the atmosphere which about two billion years ago turned 20% oxygen - oxidizing this iron to the plus three state which is virtually insoluble in water. This change from the relatively plentiful and soluble FeII, took a heavy toil on almost everything alive at the time. Surviving terrestrial and ocean-dwelling microbes developed soluble siderophore molecules to regain access to this plentiful, but otherwise inaccessible essential resource, which used hydroxamate or catechol chelating groups to bring the FeIII back into solution. Eventually higher organisms including animals, evolved. And animals used the energy of oxygen recombining with the hydrocarbons and carbohydrates in plant life to enable motion. Iron was essential to this process. But no animal, however, has been able to adequately deal, in the long run - meaning eighty year life spans - with the fact that iron is essential for the conversion of solar energy to movement, but is virtually insoluble in water at neutral pH, and, even worse, is toxic. Carbon, sulfur, nitrogen. calcium, magnesium, sodium, maybe ten other elements are also involved in life, but none of them have the power of iron to move electrons around, and none of them have the power to totally destroy the whole system. Iron does. Systems have evolved to maintain iron in specific useful and safe configurations - enzymes which utilize its catalytic powers, or transferrins and haemosiderins, which move it around and store it. But these are not perfect. Sometimes iron atoms are misplaced, and there are no known systems to recapture iron that has precipitated inside of a cell. In some tissues, cells overloaded with iron can be recycled or destroyed - but this doesn't work for neurons. Neurons sprout thousands of processes during their existence - reaching out to form networks of connections to other neurons. During development of the adult human brain a large percentage of cells are completely eliminated, and some new ones are added. It is a learning process. But once an area of the brain is up and running, there is nothing that can be done biologically, if a large number of its cells stop working for any reason. And the slow creep of precipitating iron over many decades is perhaps most often that reason. In less sophisticated tissues, like the liver, new stem cells can be activated, but in the brain, trained, structurally complex, interconnected neurons are needed, with thousands of projections that are accumulated over a lifetime of learning. So the result is slowly progressive neurodegenerative disease, like Parkinson's and Alzheimer's. This same basic mechanism can result in a variety of diseases. There are twenty or thirty proteins that that deal with iron in the brain - holding iron and passing it from place to place. Every new individual endowed with a new set of chromosomes is endowed with a new set of these proteins. Some combinations will be better than others and some will be dangerous individually and collectively. A mutation in a gene that codes for one of these proteins could disrupt its function - allowing iron atoms to become lost. These atoms that have been lost from the chemical groups that hold them will not always be safely returned to some structure like transferrin or haemoferritin. Some of them will react with water and be lost forever. Only they aren't really lost. They are piling up in the unlucky cell types that were the designated locations for expression of the most iron-leaky proteins. And oxides of iron are not just taking up critical space. Iron is very reactive. The infamous "Reactive Oxygen Species" which have been suspected of causing so many age related illnesses may just derive from various forms of iron. It is time for specialists trained in chemistry, and with an eye to the chemistry of iron, to pay some attention to neurodegenerative disease. Chris Smith Kary Mullis telling the story of iron, the element that we can't do without, but which at the same time could hold the key to our neurological downfall. Next time on Chemistry in its Element Johnny Ball will tell the story of Marie Curie and the element that she discovered and then named after her homeland. Johnny Ball Pitchblende, a uranium bearing ore, seemed to be far too radio active than could be accounted for by the uranium. They sieved and sorted by hand ounce by ounce through tons of pitchblende in a drafty, freezing shed, before eventually tiny amounts of polonium were discovered. Chris Smith So be radioactive or at least podcast proactive and join us for the mysterious story of Polonium on next week's Chemistry in its Element. I'm Chris Smith, thank you for listening, see you next time. (Promo) Chemistry in its element is brought to you by the Royal Society of Chemistry and produced by thenakedscientists.com. There's more information and other episodes of Chemistry in its element on our website at chemistryworld.org/elements. (End promo)
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Visual Elements images and videos
© Murray Robertson 1998-2017.
W. M. Haynes, ed., CRC Handbook of Chemistry and Physics, CRC Press/Taylor and Francis, Boca Raton, FL, 95th Edition, Internet Version 2015, accessed December 2014.
Tables of Physical & Chemical Constants, Kaye & Laby Online, 16th edition, 1995. Version 1.0 (2005), accessed December 2014.
J. S. Coursey, D. J. Schwab, J. J. Tsai, and R. A. Dragoset, Atomic Weights and Isotopic Compositions (version 4.1), 2015, National Institute of Standards and Technology, Gaithersburg, MD, accessed November 2016.
T. L. Cottrell, The Strengths of Chemical Bonds, Butterworth, London, 1954.
John Emsley, Nature’s Building Blocks: An A-Z Guide to the Elements, Oxford University Press, New York, 2nd Edition, 2011.
Thomas Jefferson National Accelerator Facility - Office of Science Education, It’s Elemental - The Periodic Table of Elements, accessed December 2014.
Periodic Table of Videos, accessed December 2014.
Derived in part from material provided by the British Geological Survey © NERC.
Elements 1-112, 114, 116 and 117 © John Emsley 2012. Elements 113, 115, 117 and 118 © Royal Society of Chemistry 2017.
Produced by The Naked Scientists.
Created by video journalist Brady Haran working with chemists at The University of Nottingham.
© Murray Robertson 1998-2017.
Data
W. M. Haynes, ed., CRC Handbook of Chemistry and Physics, CRC Press/Taylor and Francis, Boca Raton, FL, 95th Edition, Internet Version 2015, accessed December 2014.
Tables of Physical & Chemical Constants, Kaye & Laby Online, 16th edition, 1995. Version 1.0 (2005), accessed December 2014.
J. S. Coursey, D. J. Schwab, J. J. Tsai, and R. A. Dragoset, Atomic Weights and Isotopic Compositions (version 4.1), 2015, National Institute of Standards and Technology, Gaithersburg, MD, accessed November 2016.
T. L. Cottrell, The Strengths of Chemical Bonds, Butterworth, London, 1954.
Uses and properties
John Emsley, Nature’s Building Blocks: An A-Z Guide to the Elements, Oxford University Press, New York, 2nd Edition, 2011.
Thomas Jefferson National Accelerator Facility - Office of Science Education, It’s Elemental - The Periodic Table of Elements, accessed December 2014.
Periodic Table of Videos, accessed December 2014.
Supply risk data
Derived in part from material provided by the British Geological Survey © NERC.
History text
Elements 1-112, 114, 116 and 117 © John Emsley 2012. Elements 113, 115, 117 and 118 © Royal Society of Chemistry 2017.
Podcasts
Produced by The Naked Scientists.
Periodic Table of Videos
Created by video journalist Brady Haran working with chemists at The University of Nottingham.
