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This article is about the element. For other uses, see Chromium (disambiguation).
Chromium, 24Cr
Chromium
Appearancesilvery metallic
Standard atomic weight Ar°(Cr)
Chromium in the periodic table
Hydrogen Helium
Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon
Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon
Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton
Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon
Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon
Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson
–
↑
Cr
↓
Mo
vanadium ← chromium → manganese
Atomic number (Z)24
Groupgroup 6
Periodperiod 4
Block  d-block
Electron configuration[Ar] 3d5 4s1
Electrons per shell2, 8, 13, 1
Physical properties
Phase at STPsolid
Melting point2180 K ​(1907 °C, ​3465 °F)
Boiling point2944 K ​(2671 °C, ​4840 °F)
Density (at 20° C)7.192 g/cm3 [3]
when liquid (at m.p.)6.3 g/cm3
Heat of fusion21.0 kJ/mol
Heat of vaporization347 kJ/mol
Molar heat capacity23.35 J/(mol·K)
Vapor pressure
P (Pa) 1 10 100 1 k 10 k 100 k
at T (K) 1656 1807 1991 2223 2530 2942
Atomic properties
Oxidation statescommon: +3, +6
−4,? −2,[4] −1,[4] 0,? +1,[4] +2,[4] +4,[4] +5[4]
ElectronegativityPauling scale: 1.66
Ionization energies
  • 1st: 652.9 kJ/mol
  • 2nd: 1590.6 kJ/mol
  • 3rd: 2987 kJ/mol
  • (more)
Atomic radiusempirical: 128 pm
Covalent radius139±5 pm
Color lines in a spectral range
Spectral lines of chromium
Other properties
Natural occurrenceprimordial
Crystal structure ​body-centered cubic (bcc) (cI2)
Lattice constant
Body-centered cubic crystal structure for chromium
a = 288.49  pm (at 20 °C)[3]
Thermal expansion4.81×10−6/K (at 20 °C)[3]
Thermal conductivity93.9 W/(m⋅K)
Electrical resistivity125 nΩ⋅m (at 20 °C)
Magnetic orderingantiferromagnetic (rather: SDW)[5]
Molar magnetic susceptibility+280.0×10−6 cm3/mol (273 K)[6]
Young's modulus279 GPa
Shear modulus115 GPa
Bulk modulus160 GPa
Speed of sound thin rod5940 m/s (at 20 °C)
Poisson ratio0.21
Mohs hardness8.5
Vickers hardness1060 MPa
Brinell hardness687–6500 MPa
CAS Number7440-47-3
History
Discovery and first isolationLouis Nicolas Vauquelin (1794, 1797)
Isotopes of chromium
Main isotopes[7] Decay
abun­dance half-life (t1/2) mode pro­duct
50Cr 4.34% stable
51Cr synth 27.7025 d ε 51V
γ –
52Cr 83.8% stable
53Cr 9.50% stable
54Cr 2.37% stable
 Category: Chromium
| references

Chromium (Template:PronEng, KROH-mee-əm) is a chemical element which has the symbol Cr and atomic number 24, first element in Group 6. It is a steely-gray, lustrous, hard metal that takes a high polish and has a high melting point. It is also odorless, tasteless, and malleable. The name of the element is derived from the Greek word "chrōma" (χρωμα), meaning color, because many of its compounds are intensely colored. It was discovered by Louis Nicolas Vauquelin in the mineral crocoite (lead chromate) in 1797. Crocoite was used as a pigment, and after the discovery that the mineral chromite also contains chromium this latter mineral was used to produce pigments as well.

Chromium was regarded with great interest because of its high corrosion resistance and hardness. A major development was the discovery that steel could be made highly resistant to corrosion and discoloration by adding chromium and nickel to form stainless steel. This application, along with chrome plating (electroplating with chromium) are currently the highest-volume uses of the metal. Chromium and ferrochromium are produced from the single commercially viable ore, chromite, by silicothermic or aluminothermic reaction or by roasting and leaching processes. Although trivalent chromium (Cr(III)) is required in trace amounts for sugar and lipid metabolism in humans and its deficiency may cause a disease called chromium deficiency, hexavalent chromium (Cr(VI)) is toxic and carcinogenic, so that abandoned chromium production sites need environmental cleanup.

Characteristics

Occurrence

Chromium is the 21st most abundant element in Earth's crust with an average concentration of 100 ppm.[8] Chromium compounds are found in the environment, due to erosion of chromium-containing rocks and can be distributed by volcanic eruptions. The concentrations range in soil is between 1 and 3000 mg/kg, in sea water 5 to 800 µg/liter, and in rivers and lakes 26 µg/liter to 5.2 mg/liter.[9] The relation between Cr(III) and Cr(VI) strongly depends on pH and oxidative properties of the location, but in most cases, the Cr(III) is the dominating species,[9] although in some areas the ground water can contain up to 39 µg of total chromium of which 30 µg is present as Cr(VI).[10]

Chromite ore

Chromium is mined as chromite (FeCr2O4) ore.[11] About two-fifths of the chromite ores and concentrates in the world are produced in South Africa, while Kazakhstan, India, Russia, and Turkey are also substantial producers. Untapped chromite deposits are plentiful, but geographically concentrated in Kazakhstan and southern Africa.[12]

Though native chromium deposits are rare, some native chromium metal has been discovered.[13][14] The Udachnaya Pipe in Russia produces samples of the native metal. This mine is a kimberlite pipe rich in diamonds, and the reducing environment helped produce both elemental chromium and diamond.[15]

Isotopes

Main article: Isotopes of chromium

Naturally occurring chromium is composed of three stable isotopes; 52Cr, 53Cr and 54Cr with 52Cr being the most abundant (83.789% natural abundance). Nineteen radioisotopes have been characterized with the most stable being 50Cr with a half-life of (more than) 1.8×1017 years, and 51Cr with a half-life of 27.7 days. All of the remaining radioactive isotopes have half-lives that are less than 24 hours and the majority of these have half-lives that are less than 1 minute. This element also has 2 meta states.[16]

53Cr is the radiogenic decay product of 53Mn. Chromium isotopic contents are typically combined with manganese isotopic contents and have found application in isotope geology. Mn-Cr isotope ratios reinforce the evidence from 26Al and 107Pd for the early history of the solar system. Variations in 53Cr/52Cr and Mn/Cr ratios from several meteorites indicate an initial 53Mn/55Mn ratio that suggests Mn-Cr isotopic composition must result from in-situ decay of 53Mn in differentiated planetary bodies. Hence 53Cr provides additional evidence for nucleosynthetic processes immediately before coalescence of the solar system.[17]

The isotopes of chromium range in atomic mass from 43 u (43Cr) to 67 u (67Cr). The primary decay mode before the most abundant stable isotope, 52Cr, is electron capture and the primary mode after is beta decay.[16] 53Cr has been posited as a proxy for atmospheric oxygen concentration.[18]

Chemistry

Oxidation states
of chromium[note 1][19]
−2 Na
2[Cr(CO)
5]
−1 Na
2[Cr
2(CO)
10]
0 Cr(C
6H
6)
2
+1 K
3[Cr(CN)
5NO]
+2 CrCl
2
+3 CrCl
3
+4 K
2CrF
6
+5 K
3CrO
8
+6 K
2CrO
4

Chromium is a member of the transition metals, in group 6. Chromium(0) has an electronic configuration of 4s13d5, due to the lower energy of the high spin configuration. Chromium exhibits a wide range of possible oxidation states. The most common oxidation states of chromium are +2, +3, and +6, with +3 being the most stable. The +1, +4 and +5 states are rare.

The following is the Pourbaix diagram for chromium in pure water, perchloric acid or sodium hydroxide:[9][20]

Chromium(III)

The oxidation state +3 is the most stable, and a large number of chromium(III) compounds are known. Chromium(III) can be obtained by dissolving elemental chromium in acids like hydrochloric acid or sulfuric acid. The Cr3+
 ion has a similar radius (63 pm) to the Al3+
 ion (radius 50 pm), so they can replace each other in some compounds, such as in chrome alum and alum. When a trace amount of Cr3+
 replaces Al3+
 in corundum (aluminium oxide, Al2O3), the red-colored ruby is formed.

Chromium(III) chloride hexahydrate ([CrCl2(H2O)4]Cl·2H2O)

Chromium ions tend to form complexes; chromium ions in water are usually octahedrally coordinated with water molecules to form hydrates. The commercially available chromium(III) chloride hydrate is the dark green complex [CrCl2(H2O)4]Cl, but two other forms are known: pale green [CrCl(H2O)5]Cl2, and the violet [Cr(H2O)6]Cl3. If water-free green chromium(III) chloride is dissolved in water then the green solution turns violet after some time, due to the substitution of water for chloride in the inner coordination sphere. This kind of reaction is also observed in chrome alum solutions and other water-soluble chromium(III) salts. The reverse reaction may be induced by heating the solution.

Chromium(III) hydroxide (Cr(OH)3) is amphoteric, dissolving in acidic solutions to form [Cr(H2O)6]3+, and in basic solutions to form [Cr(OH)
6]3−
. It is dehydrated by heating to form the green chromium(III) oxide (Cr2O3), which is the stable oxide with a crystal structure identical to that of corundum.[21]

Chromium(III) chloride (CrCl3)

Chromium(VI)

Chromium(VI) compounds are powerful oxidants, and, except the hexafluoride, contain oxygen as a ligand, such as the chromate anion (CrO2−
4) and chromyl chloride (CrO
2Cl
2).[21]

Chromium(VI) is most commonly encountered in the chromate (CrO2−
4) and dichromate (Cr
2O2−
7) anions. Chromate is produced industrially by the oxidative roasting of chromite ore with calcium or sodium carbonate. The chromate and dichromate anions are in equilibrium:

2 CrO2−
4 + 2 H
3O+
 → Cr
2O2−
7 + 3 H
2O

The dominant species is therefore, by the law of mass action, determined by the pH of the solution. The change in equilibrium is visible by a change from yellow (chromate) to orange (dichromate), such as when an acid is added to a neutral solution of potassium chromate. At yet lower pH values, further condensation to more complex oxyanions of chromium is possible.

Both the chromate and dichromate anions are strong oxidizing reagents at low pH:[21]

Cr
2O2−
7 + 14 H
3O+
 + 6 e− → 2 Cr3+
 + 21 H
2O (ε0 = 1.33 V)

However, they are only moderately oxidizing at high pH:[21]

CrO2−
4 + 4 H
2O + 3 e− → Cr(OH)3+
 + 5 OH−
 (ε0 = −0.13 V)
Chromium(VI) oxide

Chromium(VI) compounds in solution can be detected by adding an acidic hydrogen peroxide solution. The unstable dark blue chromium(VI) peroxide (CrO5) is formed, which can be stabilized as an ether adduct CrO
5·OR
2.[21]

Chromic acid has the hypothetical structure H
2CrO
4. Neither chromic nor dichromic acid can be isolated, but their anions are found in a variety of compounds, the chromates and dichromates. The dark red chromium(VI) oxide CrO
3, the acid anhydride of chromic acid, is sold industrially as "chromic acid".[21] It can be produced by mixing sulfuric acid with dichromate, and is an extremely strong oxidizing agent.

Sodium chromate

Other oxidation states

The oxidation state +5 is only realized in few compounds. The only binary compound is the highly volatile chromium(V) fluoride (CrF5). This red solid has a melting point of 30°C and a boiling point of 117°C, and can be synthesized by reacting fluorine with chromium at 400°C and 200 bar pressure. The peroxochromate(V) is another example of the +5 oxidation state. Potassium peroxochromate (K3[Cr(O2)4]) is made by reacting potassium chromate with hydrogen peroxide at low temperatures. This red brown compound is stable at room temperature but decomposes spontaneously at 150–170 °C.[22]

Chromium(IV) compounds (in the +4 oxidation state) are slightly more stable than the chromium(V) compounds. The tetrahalides, CrF4, CrCl4, and CrBr4, can be produced by reacting the trihalides (CrX
3) with excess amounts of the corresponding halogen at elevated temperatures. Most of the compounds are susceptible to disproportionation reactions and are not stable in water.

An example of a stable chromium(II) compound is the water-stable chromium(II) chloride, CrCl
2, which can be made by reduction of chromium(III) chloride with zinc. The resulting bright blue solution is only stable at neutral pH when the solution is very pure.[21]

Passivation

Chromium metal left standing in air is passivated by oxygen, forming a thin protective oxide surface layer. This layer is a spinel structure only a few atoms thick. It is very dense, and prevents the diffusion of oxygen into the underlying material. This is in contrast to iron or plain carbon steels, where the oxygen migrates into the underlying material and causes rusting.[23] The passivation can be increased by short contact with oxidizing acids like nitric acid. Passivated chromium is stable against acids. The opposite effect can be achieved by treatment with a strong reducing reactant that destroys the protective oxide layer on the metal. Chromium metal treated in this way readily dissolves in weak acids.[21]

Chromium, unlike metals such as iron and nickel, does not suffer from hydrogen embrittlement. However, it does suffer from nitrogen embrittlement, reacting with nitrogen from air and forming brittle nitrides at the high temperatures necessary to work the metal parts.[24]

Quintuple bond

Chromium compound determined experimentally to contain a Cr-Cr quintuple bond

Chromium is notable for its ability to form quintuple covalent bonds. The product of a reaction between chromium(I) and a hydrocarbon radical was shown via X-ray diffraction to contain a quintuple bond of length 183.51(4) pm joining the two central chromium atoms.[25] Extremely bulky monodentate ligands stabilize this compound by shielding the quintuple bond from further reactions.

Physical properties

Chromium is remarkable for its magnetic properties: it is the only elemental solid which shows antiferromagnetic ordering at room temperature (and below). Above 38 °C, it transforms into a paramagnetic state.[5]

History

Crocoite (PbCrO4)

Weapons found in burial pits dating from the late 3rd century BC Qin Dynasty of the Terracotta Army near Xi'an, China have been analyzed by archaeologists. Although buried more than 2,000 years ago, the ancient bronze tips of crossbow bolts and swords found at the site showed no sign of corrosion, because the bronze was coated with chromium.[26]

Chromium came to the attention of westerners in the 18th century. On 26 July 1761, Johann Gottlob Lehmann found an orange-red mineral in the Beryozovskoye mines in the Ural Mountains which he named Siberian red lead. Though misidentified as a lead compound with selenium and iron components, the mineral was Crocoite (lead chromate) with a formula of PbCrO4.[27]

In 1770, Peter Simon Pallas visited the same site as Lehmann and found a red lead mineral that had useful properties as a pigment in paints. The use of Siberian red lead as a paint pigment developed rapidly. A bright yellow pigment made from crocoite also became fashionable.[27]

Ruby is colored by a small amount of chromium

In 1797, Louis Nicolas Vauquelin received samples of crocoite ore. He produced chromium oxide (CrO3) by mixing crocoite with hydrochloric acid. In 1798, Vauquelin discovered that he could isolate metallic chromium by heating the oxide in a charcoal oven.[28] He was also able to detect traces of chromium in precious gemstones, such as ruby or emerald.[27][29]

During the 1800s, chromium was primarily used as a component of paints and in tanning salts. At first, crocoite from Russia was the main source, but in 1827, a larger chromite deposit was discovered near Baltimore, United States. This made the United states the largest producer of chromium products till 1848 when large deposits of chromite where found near Bursa, Turkey.[11]

Chromium is also known for its luster when polished. It is used as a protective and decorative coating on car parts, plumbing fixtures, furniture parts and many other items, usually applied by electroplating. Chromium was used for electroplating as early as 1848, but this use only became widespread with the development of an improved process in 1924.[30]

Metal alloys now account for 85% of the use of chromium. The remainder is used in the chemical industry and refractory and foundry industries.

Production

World production trend of chromium

Approximately 4.4 million metric tons of marketable chromite ore were produced in 2000, and converted into ~3.3 million tons of ferro-chrome with an approximate market value of 2.5 billion United States dollars.[31] The largest producers of chromium ore have been South Africa (44%) India (18%), Kazakhstan (16%) Zimbabwe (5%), Finland (4%) Iran (4%) and Brazil (2%) with several other countries producing the rest of less than 10% of the world production.[31]

The two main products of chromium ore refining are ferrochromium and metallic chromium. For those products the ore smelter process differs considerably. For the production of ferrochromium, the chromite ore (FeCr2O4) is reduced in large scale in electric arc furnace or in smaller smelters with either aluminium or silicon in an aluminothermic reaction.[32]

Chromium ore output in 2002[31]

For the production of pure chromium, the iron has to be separated from the chromium in a two step roasting and leaching process. The chromite ore is heated with a mixture of calcium carbonate and sodium carbonate in the presence of air. The chromium is oxidized to the hexavalent form, while the iron forms the stable Fe2O3. The subsequent leaching at higher elevated temperatures dissolves the chromates and leaves the insoluble iron oxide. The chromate is converted by sulfuric acid into the dichromate.[32]

4 FeCr2O4 + 8 Na2CO3 + 7 O2 → 8 Na2CrO4 + 2 Fe2O3 + 8 CO2
2 Na2CrO4 + H2SO4 → Na2Cr2O7 + Na2SO4 + H2O

The dichromate is converted to the chromium(III) oxide by reduction with carbon and then reduced in an aluminothermic reaction to chromium.[32]

Na2Cr2O7 + 2 C → Cr2O3 + Na2CO3 + CO
Cr2O3 + 2 Al → Al2O3 + 2 Cr

Applications

Metallurgy

Decorative chrome plating on a motorcycle.

The strengthening effect of forming stable metal carbides at the grain boundaries and the strong increase in corrosion resistance made chromium an important alloying material for steel. The high speed tool steels contain between 3 and 5% chromium. An important stainless steel is 18/10 stainless, made from iron with 10% nickel and 18% chromium, is widely used for cookware and cutlery. For these applications, ferrochromium is added to the molten iron. Also nickel-based alloys increase in strength due to the formation of discrete, stable metal carbide particles at the grain boundaries. For example, Inconel 718 contains 18.6% chromium. Because of the excellent high temperature properties of these nickel superalloys, they are used in jet engines and gas turbines in lieu of common structural materials.[33]

The relative high hardness and corrosion resistance of unalloyed chromium makes it a good surface coating. A thin layer of chromium is deposited on pretreated metallic surfaces by electroplating techniques. There are two deposition methods: Thin, below 1 µm thickness, layers are deposited by chrome plating, and are used for decorative surfaces. If wear-resistant surfaces are needed then thicker chromium layers of up to mm thickness are deposited. Both methods normally use acidic chromate or dichromate solutions. To prevent the energy consuming change in oxidation state, the use of Chromium(III) sulfate is under development, but for most applications, the established process is used.[30]

In the chromate conversion coating process, the strong oxidative properties of chromates are used to deposit a protective oxide layer on metals like aluminium, zinc and cadmium. This passivation and the self healing properties by the chromate stored in the chromate conversion coating, which is capable to migrate to local defects, are the benefits of this coating method.[34] Because of environmental and health regulations on chromates, alternative coating method are under development.[35]

Anodizing of aluminium is another electrochemical process, which does not lead to the deposition of chromium, but uses chromic acid as electrolyte in the solution. During anodization, an oxide layer is formed on the aluminium. The use of chromic acid, instead of the normally used sulfuric acid, leads to a slight difference of these oxide layers.[36] The high toxicity of Cr(VI) compounds, used in the established chromium electroplating process, and the strengthening of safety and environmental regulations demand a search for substitutes for chromium or at least a change to less toxic chromium(III) compounds.[30]

Dye and pigment

School bus painted in Chrome yellow[37]

The mineral crocoite (lead chromate PbCrO4) was used as a yellow pigment shortly after its discovery. After a synthesis method became available starting from the more abundant chromite, Chrome yellow was, together with cadmium yellow, one of the most used yellow pigments. The pigment does not degrade in the light and has a strong color. The signaling effect of yellow was used for school buses in the United States and for Postal Service (for example Deutsche Post) in Europe. The use of chrome yellow declined due to environmental and safety concerns and was substituted by organic pigments or other lead-free alternatives.[38] Other pigments based on chromium are, for example, the bright red pigment Chrome red, which is a basic lead chromate (PbCrO4•Pb(OH)2).[38] Chrome green is a mixture of Prussian blue and chrome yellow, while the Chrome oxide green is Chromium(III) oxide.[38]

Glass is colored green by the addition of chromium(III) oxide. This is similar to emerald, which is also colored by chromium.[39] A red color is achieved by doping chromium(III) into the crystals of corundum, which are then called ruby. Therefore, chromium is used in producing synthetic rubies.[40]

The toxicity of chromium(VI) salts is used in the preservation of wood. For example, chromated copper arsenate (CCA) is used in timber treatment to prevent wood from decay fungi, wood attacking insects, including termites, and marine borers.[41] The formulations contain chromium based on the oxide CrO3 between 35.3% and 65.5%. In the United States, 65,300 metric tons of CCA solution have been used in 1996.[41]

Tanning

Chromium(III) salts, especially chrome alum and chromium(III) sulfate, are used in the tanning of leather. The chromium(III) stabilizes the leather by cross linking the collagen fibers within the leather.[42] Chromium tanned leather can contain between 4 and 5% of chromium, which is tightly bound to the proteins.[11]

Refractory material

The high heat resistivity and high melting point makes chromite and chromium(III) oxide a material for high temperature refractory applications, like blast furnaces, cement kilns, molds for the firing of bricks and as foundry sands for the casting of metals. In these applications, the refractory materials are made from mixtures of chromite and magnesite. The use is declining because of the environmental regulations due to the possibility of the formation of chromium(VI).[32]

Other use

Several chromium compounds are used as catalyst. For example the Phillips catalysts for the production polyethylene are mixtures of chromium and silicon dioxide or mixtures of chromium and titanium and aluminium oxide.[43] Chromium(IV) oxide (CrO2) is a magnetic compound. Its ideal shape anisotropy, which imparted high coercivity and remanent magnetization, made it a compound superior to the γ-Fe2O3. Chromium(IV) oxide is used to manufacture magnetic tape used in high performance audio tape and standard audio cassette.[44] Chromates can prevent corrosion of steel under wet conditions, and therefore chromates are added to the drilling muds.[45] The long known influence of chromium uptake on diabetes conditions suggested the positive influence of dietary supplement containing chromium(III) also on healthy persons. For this reason, dietary supplement or slimming aid usually contain chromium(III) chloride, chromium(III) picolinate, chromium(III) polynicotinate or amino acid chelate, such as chromium(III) D-phenylalanine. The benefit of those supplements is still under investigation and is questioned by some studies.[46][47]

  • Chromium hexacarbonyl Cr(CO)6 is used as a gasoline additive.[48]
  • Chromium(III) oxide is a metal polish known as green rouge.
  • Chromic acid is a powerful oxidizing agent and is a useful compound for cleaning laboratory glassware of any trace of organic compounds. It is prepared in situ by dissolving potassium dichromate in concentrated sulfuric acid, which is then used to wash the apparatus. Sodium dichromate is sometimes used because of its higher solubility (5 g/100 ml vs. 20 g/100 ml respectively). Potassium dichromate is a chemical reagent, used in cleaning laboratory glassware and as a titrating agent. It is also used as a mordant (i.e., a fixing agent) for dyes in fabric.

Biological role

Trivalent chromium (Cr(III) or Cr3+) in trace amounts influences sugar and lipid metabolism in humans, and its deficiency is suspected to cause a disease called chromium deficiency.[49] In contrast, hexavalent chromium (Cr(VI) or Cr6+) is very toxic and mutagenic when inhaled. Cr(VI) has not been established as a carcinogen when in solution, though it may cause allergic contact dermatitis (ACD).[50]

The use of chromium-containing dietary supplements is controversial due to the complex effects of the used supplements.[51] The popular dietary supplement chromium picolinate complex generates chromosome damage in hamster cells.[52] In the United States the dietary guidelines for daily chromium uptake were lowered from 50-200 µg for an adult to 35 µg (adult male) and to 25 µg (adult female).[53]

Precautions

Water insoluble chromium(III) compounds and chromium metal are not considered a health hazard, while the toxicity and carcinogenic properties of chromium(VI) have been known for a long time.[54] An actual investigation into hexavalent chromium release into drinking water was used as the plot-basis of the motion picture Erin Brockovich.

Because of the specific transport mechanisms, only limited amounts of chromium(III) enter the cells. Several in vitro studies indicated that high concentrations of chromium(III) in the cell can lead to DNA damage.[55] Acute oral toxicity ranges between 1500 and 3300 µg/kg.[56] The proposed beneficial effects of chromium(III) and the use as dietary supplements yielded some controversial results, but recent reviews suggest that moderate uptake of chromium(III) through dietary supplements poses no risk.[55]

World Health Organization recommended maximum allowable concentration in drinking water for chromium (VI) is 0.05 milligrams per liter. Hexavalent chromium is also one of the substances whose use is restricted by the European Restriction of Hazardous Substances Directive.

The acute oral toxicity for chromium(VI) ranges between 50 and 150 µg/kg.[56] In the body, chromium(VI) is reduced by several mechanisms to chromium(III) already in the blood before it enters the cells. The chromium(III) is excreted from the body, whereas the chromate ion is transferred into the cell by a transport mechanism, by which also sulfate and phosphate ions enter the cell. The acute toxicity of chromium(VI) is due to its strong oxidational properties. After it reaches the blood stream, it damages the kidneys, the liver and blood cells through oxidation reactions. Hemolysis, renal and liver failure are the results of these damages. Aggressive dialysis can improve the situation.[57]

The carcinogenity of chromate dust is known for a long time, and in 1890 the first publication described the elevated cancer risk of workers in a chromate dye company.[58][59] Three mechanisms have been proposed to describe the genotoxicity of chromium(VI). The first mechanism includes highly reactive hydroxyl radicals and other reactive radicals which are by products of the reduction of chromium(VI) to chromium(III). The second process includes the direct binding of chromium(V), produced by reduction in the cell, and chromium(IV) compounds to the DNA. The last mechanism attributed the genotoxicity to the binding to the DNA of the end product of the chromium(III) reduction.[60]

Chromium salts (chromates) are also the cause of allergic reactions in some people. Chromates are often used to manufacture, amongst other things, leather products, paints, cement, mortar and anti-corrosives. Contact with products containing chromates can lead to allergic contact dermatitis and irritant dermatitis, resulting in ulceration of the skin, sometimes referred to as "chrome ulcers". This condition is often found in workers that have been exposed to strong chromate solutions in electroplating, tanning and chrome-producing manufacturers.[61][62][63]

In some parts of Russia, pentavalent chromium was reported as one of the causes of premature dementia.[64]

Environmental issues

As chromium compounds were used in dyes and paints and the tanning of leather, these compounds are often found in soil and groundwater at abandoned industrial sites, now needing environmental cleanup and remediation per the treatment of brownfield land. Primer paint containing hexavalent chromium is still widely used for aerospace and automobile refinishing applications.

See also

Notes

  1. ^ Common oxidation states are in bold.

References

  1. ^ "Standard Atomic Weights: Chromium". CIAAW. 1983.
  2. ^ Prohaska, Thomas; Irrgeher, Johanna; Benefield, Jacqueline; Böhlke, John K.; Chesson, Lesley A.; Coplen, Tyler B.; Ding, Tiping; Dunn, Philip J. H.; Gröning, Manfred; Holden, Norman E.; Meijer, Harro A. J. (2022-05-04). "Standard atomic weights of the elements 2021 (IUPAC Technical Report)". Pure and Applied Chemistry. doi:10.1515/pac-2019-0603. ISSN 1365-3075.
  3. ^ a b c Arblaster, John W. (2018). Selected Values of the Crystallographic Properties of Elements. Materials Park, Ohio: ASM International. ISBN 978-1-62708-155-9.
  4. ^ a b c d e f Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. p. 28. ISBN 978-0-08-037941-8.
  5. ^ a b Fawcett, Eric (1988). "Spin-density-wave antiferromagnetism in chromium". Reviews of Modern Physics. 60: 209. Bibcode:1988RvMP...60..209F. doi:10.1103/RevModPhys.60.209.
  6. ^ Weast, Robert (1984). CRC, Handbook of Chemistry and Physics. Boca Raton, Florida: Chemical Rubber Company Publishing. pp. E110. ISBN 0-8493-0464-4.
  7. ^ Kondev, F. G.; Wang, M.; Huang, W. J.; Naimi, S.; Audi, G. (2021). "The NUBASE2020 evaluation of nuclear properties" (PDF). Chinese Physics C. 45 (3): 030001. doi:10.1088/1674-1137/abddae.
  8. ^ Emsley, John (2001). "Chromium". Nature's Building Blocks: An A-Z Guide to the Elements. Oxford, England, UK: Oxford University Press. pp. 495–498. ISBN 0198503407.
  9. ^ a b c Kotaś, J.; Stasicka, Z (2000). "Chromium occurrence in the environment and methods of its speciation". Environmental Pollution. 107 (3): 263–283. doi:10.1016/S0269-7491(99)00168-2. PMID 15092973.
  10. ^ Gonzalez, A. R.; Ndung'u, K; Flegal, AR (2005). "Natural Occurrence of Hexavalent Chromium in the Aromas Red Sands Aquifer, California". Environmental Science and Technology. 39 (15): 5505–5511. doi:10.1021/es048835n. PMID 16124280.
  11. ^ a b c National Research Council (U.S.). Committee on Biologic Effects of Atmospheric Pollutants (1974). Chromium. National Academy of Sciences. p. 155. ISBN 9780309022170.
  12. ^ Papp, John F. "Commodity Summary 2009: Chromium" (PDF). United States Geological Survey. Retrieved 2009-03-17.
  13. ^ Fleischer, Michael (l982). "New Mineral Names" (PDF). American Mineralogist. 67: 854–860. {{cite journal}}: Check date values in: |year= (help)CS1 maint: year (link)
  14. ^ http://www.mindat.org/min-1037.html Mindat with location data
  15. ^ http://www.mindat.org/locentry-27628.html Mindat
  16. ^ a b Georges, Audi (2003). "The NUBASE Evaluation of Nuclear and Decay Properties". Nuclear Physics A. 729. Atomic Mass Data Center: 3–128. doi:10.1016/j.nuclphysa.2003.11.001.
  17. ^ Birck, J. L.; Rotaru, M; Allegre, C (1999). "53Mn-53Cr evolution of the early solar system". Geochimica et Cosmochimica Acta. 63 (23–24): 4111–4117. doi:10.1016/S0016-7037(99)00312-9.
  18. ^ Attention: This template ({{cite doi}}) is deprecated. To cite the publication identified by doi:10.1038/nature08266, please use {{cite journal}} (if it was published in a bona fide academic journal, otherwise {{cite report}} with |doi=10.1038/nature08266 instead.
  19. ^ Schmidt, Max (1968). "VI. Nebengruppe". Anorganische Chemie II (in German). Wissenschaftsverlag. pp. 119–127.
  20. ^ Ignasi Puigdomenech, Hydra/Medusa Chemical Equilibrium Database and Plotting Software (2004) KTH Royal Institute of Technology, freely downloadable software at [1]
  21. ^ a b c d e f g h Holleman, Arnold F. (1985). "Chromium". Lehrbuch der Anorganischen Chemie (in German) (91–100 ed.). Walter de Gruyter. pp. 1081–1095. ISBN 3110075113. {{cite book}}: Unknown parameter |coauthors= ignored (|author= suggested) (help)CS1 maint: extra punctuation (link)
  22. ^ Haxhillazi, Gentiana. "Preparation, Structure and Vibrational Spectroscopy of Tetraperoxo Complexes of CrV+, VV+, NbV+ and TaV+ (PhD thesis, University of Siegen, 2003)". {{cite web}}: Unknown parameter |publisher >= ignored (help)
  23. ^ Wallwork, G. R. (1976). "The oxidation of alloys" (PDF). Reports on the Progress Physics. 39: 401–485. doi:10.1088/0034-4885/39/5/001.
  24. ^ National Research Council (U.S.). Committee on Coatings (1970). High-temperature oxidation-resistant coatings: coatings for protection from oxidation of superalloys, refractory metals, and graphite. National Academy of Sciences. ISBN 0309017696.
  25. ^ T. Nguyen, A. D. Sutton, M. Brynda, J. C. Fettinger, G. J. Long and P. P. Power (2005). "Synthesis of a Stable Compound with Fivefold Bonding Between Two Chromium(I) Centers". Science. 310 (5749): 844–847. doi:10.1126/science.1116789. PMID 16179432.{{cite journal}}: CS1 maint: multiple names: authors list (link)
  26. ^ Cotterell, Maurice. (2004). The Terracotta Warriors: The Secret Codes of the Emperor's Army. Rochester: Bear and Company. ISBN 159143033X. Page 102.
  27. ^ a b c Jacques Guertin, James Alan Jacobs, Cynthia P. Avakian, (2005). Chromium (VI) Handbook. CRC Press. pp. 7–11. ISBN 9781566706087.{{cite book}}: CS1 maint: extra punctuation (link) CS1 maint: multiple names: authors list (link)
  28. ^ Vauquelin, Louis Nicolas (1798). "Memoir on a New Metallic Acid which exists in the Red Lead of Sibiria". Journal of Natural Philosophy, Chemistry, and the Art. 3: 146.
  29. ^ van der Krogt, Peter. "Chromium". Retrieved 2008-08-24.
  30. ^ a b c Dennis, J. K.; Such, T. E. (1993). "History of Chromium Plating". Nickel and Chromium Plating. Woodhead Publishing. pp. 9–12. ISBN 9781855730816.{{cite book}}: CS1 maint: multiple names: authors list (link)
  31. ^ a b c Papp, John F. "Mineral Yearbook 2002: Chromium" (PDF). United States Geological Survey. Retrieved 2009-02-16.
  32. ^ a b c d Papp, John F (2006). "Chromite". Industrial Minerals & Rocks: Commodities, Markets, and Uses (7th ed.). SME. ISBN 9780873352338. {{cite book}}: Unknown parameter |coauthor= ignored (|author= suggested) (help)
  33. ^ Bhadeshia, H. K. D. H. "Nickel-Based Superalloys". University of Cambridge. Retrieved 2009-02-17.
  34. ^ Edwards, Joseph (1997). Coating and Surface Treatment Systems for Metals. Finishing Publications Ltd. and ASM International. pp. 66–71. ISBN 0-904477-16-9.
  35. ^ Zhao, J. (2001). "Effects of chromate and chromate conversion coatings on corrosion of aluminum alloy 2024-T3" (PDF). Surface and Coatings Technology. 140 (1): 51–57. doi:10.1016/S0257-8972(01)01003-9.
  36. ^ Sprague, J. A. (1994). ASM Handbook: Surface Engineering. ASM International. ISBN 9780871703842. Retrieved 2009-02-17. {{cite book}}: Unknown parameter |coauthor= ignored (|author= suggested) (help)
  37. ^ Worobec, Mary Devine (1992). Toxic Substances Controls Guide: Federal Regulation of Chemicals in the Environment. Washington, D.C.: Bureau of National Affairs. p. 13. ISBN 9780871797520. {{cite book}}: Cite has empty unknown parameter: |unused_data= (help); Text "BNA Books" ignored (help)
  38. ^ a b c Gettens, Rutherford John (1966). "Painting Materials: A Short Encyclopaedia". Courier Dover Publications: 105-–106. ISBN 9780486215976. {{cite journal}}: Cite has empty unknown parameter: |unused_data= (help); Cite journal requires |journal= (help); Text "chapter Chrome yellow" ignored (help)
  39. ^ Carstens, Harald (1973). "The red-green change in chromium-bearing garnets". Contributions to Mineralogy and Petrology. 41 (3): 273–276. doi:10.1007/BF00371036.
  40. ^ Moss, S. C. (1964). "The chromium position in ruby" (PDF). Zeitschrift fur Kristallographie. 120: 359–363.
  41. ^ a b Hingston, J, J. A.; Collins, CD; Murphy, RJ; Lester, JN (2001). "Leaching of chromated copper arsenate wood preservatives: a review". Environmental Pollution. 111 (1): 53–66. doi:10.1016/S0269-7491(00)00030-0. PMID 11202715.
  42. ^ Brown, E. M. (1997). "A Conformational Study of Collagen as Affected by Tanning Procedures". Journal of the American Leather Chemists Association. 92: 225–233.
  43. ^ Weckhuysen, Bert M. (1999). "Olefin polymerization over supported chromium oxide catalysts". Catalysis Today. 51 (2): 215–221. doi:10.1016/S0920-5861(99)00046-2.
  44. ^ Mallinson, John C. (1993). "Chromium Dioxide". The foundations of magnetic recording. Academic Press. ISBN 9780124666269.
  45. ^ Garverick, Linda (1994). Corrosion in the Petrochemical Industry. ASM International. ISBN 9780871705051.
  46. ^ Heimbach, J.T. (2005). "Chromium: Recent Studies Regarding Nutritional Roles and Safety". Nutrition Today. 40 (4): 189–-195.
  47. ^ Vincent,, John B . (2003). "The Potential Value and Toxicity of Chromium Picolinate as a Nutritional Supplement, Weight Loss Agent and Muscle Development Agent". Sports Medicine:Volume. 33 (3): 213–230. doi:10.2165/00007256-200333030-00004.{{cite journal}}: CS1 maint: extra punctuation (link)
  48. ^ Patnaik, Pradyot (2003). "Chromium hexacarbonyl". Handbook of Inorganic Chemicals. McGraw-Hill Professional. pp. 222–223. ISBN 9780070494398.
  49. ^ Mertz, Walter (1 April 1993). "Chromium in Human Nutrition: A Review". Journal of Nutrition. 123 (4): 626–636. PMID 8463863.
  50. ^ "ToxFAQs: Chromium". Agency for Toxic Substances & Disease Registry, Centers for Disease Control and Prevention. 2001. Retrieved 2007-10-02. {{cite web}}: Unknown parameter |month= ignored (help)
  51. ^ Cronin, Joseph R. (2004). "The Chromium Controversy". Alternative and Complementary Therapies. 10 (1): 39–42. doi:10.1089/107628004772830393.
  52. ^ Stearns, D. M.; W; P; W (1 December 1995). "Chromium(III) picolinate produces chromosome damage in Chinese hamster ovary cells". Federation of American Societies for Experimental Biology. 9 (15): 1643–1648. PMID 8529845.
  53. ^ Vincent, J. B. (2007). "Recent advances in the nutritional biochemistry of trivalent chromium". Proceedings of the Nutrition Society. 63 (01): 41–47. doi:10.1079/PNS2003315.
  54. ^ Barceloux, Donald G.; Barceloux, Donald (1999). "Chromium". Clinical Toxicology. 37 (2): 173–194. doi:10.1081/CLT-100102418. PMID 10382554.
  55. ^ a b Eastmond, David A.; MacGregor, JT; Slesinski, RS (2008). "Trivalent Chromium: Assessing the Genotoxic Risk of an Essential Trace Element and Widely Used Human and Animal Nutritional Supplement". Critical Reviews in Toxicology. 38 (3): 173–190. doi:10.1080/10408440701845401. PMID 18324515.
  56. ^ a b Katz, Sidney A.; Salem, H (1992). "The toxicology of chromium with respect to its chemical speciation: A review". Journal of Applied Toxicology. 13 (3): 217–224. doi:10.1002/jat.2550130314. PMID 8326093.
  57. ^ Dayan, A. D.; Paine, AJ (2001). "Mechanisms of chromium toxicity, carcinogenicity and allergenicity: Review of the literature from 1985 to 2000". Human & Experimental Toxicology. 20 (9): 439–451. doi:10.1191/096032701682693062. PMID 11776406.
  58. ^ Newman, D. (1890). "A case of adeno-carcinoma of the left inferior turbinated body, and perforation of the nasal septum, in the person of a worker in chrome pigments". Glasgow Med J. 33: 469–470. {{cite journal}}: line feed character in |title= at position 87 (help)
  59. ^ Langard, Sverre (1990). "One Hundred Years of Chromium and Cancer: A Review of Epidemiological Evidence and Selected Case Reports". American Journal of Industrial Medicine. 17 (2): 189–215. doi:10.1002/ajim.4700170205. PMID 2405656. {{cite journal}}: line feed character in |title= at position 92 (help)
  60. ^ M. D., Cohen; Kargacin, B; Klein, CB; Costa, M (1993). "Mechanisms of chromium carcinogenicity and toxicity". Critical reviews in toxicology. 23 (3): 255–81. doi:10.3109/10408449309105012. PMID 8260068.
  61. ^ "Chrome Contact Allergy". DermNet NZ.
  62. ^ Basketter, David; Horev, L; Slodovnik, D; Merimes, S; Trattner, A; Ingber, A (2000). "Investigation of the threshold for allergic reactivity to chromium". Contact Dermatitis. 44 (2): 70–74. doi:10.1034/j.1600-0536.2001.440202.x. PMID 11205406.
  63. ^ Basketter, D. A.; Briatico-Vangosa, G; Kaestner, W; Lally, C; Bontinck, WJ (1992). "Nickel, cobalt and chromium in consumer products: a role in allergic contact dermatitis?". Contact Dermatitis. 28 (1): 15–25. doi:10.1111/j.1600-0536.1993.tb03318.x. PMID 8428439.
  64. ^ Chromium Toxicity on the Corrosion Doctors Web site maintained by Canadian Physical Chemist, Pierre R. Roberge, PhD, P.Eng. (access date 27 april 2009)
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