Cobalt

Cobalt
Atomic number	27
Symbol	Co
Group	9
Boiling point	2927 °C
Electron configuration	[Ar] 3d7 4s2
Melting point	1495 °C
Main isotopes	59Co, 60Co
Block	d
Phase STP	Solid
Oxidation states	+2, +3
Wikidata	Q740

Overview and Key Facts. Cobalt (symbol Co, atomic number 27) is a transition metal found in group 9 of the periodic table, period 4. At standard conditions it is a solid metal with a lustrous, silver-gray appearance. It is ferromagnetic at room temperature (like iron and nickel) and retains magnetization up to its Curie point (~1390 K). Cobalt atoms have 27 electrons (atomic weight ~58.93 u) and an electron configuration [Ar] 3d^7 4s^2, giving common oxidation states +2 and +3. Cobalt is relatively uncommon in Earth’s crust (~30 ppm) and occurs only in combined form (it is not found as a free metal naturally). Its most important applications arise from its high melting point, magnetic strength, and the vibrant colors of its compounds.

Symbol: Co
Atomic number: 27
Atomic weight: 58.933 (≈59)
Group/Period/Block: 9 / 4 / d-block (transition metal)
Phase at STP: solid (silvery-gray metal)
Common oxidation states: +2, +3
Electron configuration: [Ar] 3d^7 4s^2
Appearance: lustrous, hard, silvery-gray metal (brittle when unalloyed)

Atomic Structure and Electron Configuration. Cobalt atoms have 27 electrons arranged in shells of 2, 8, 15, 2. The valence electrons occupy the 3d and 4s orbitals Ar] 3d^7 4s^2) These outer electrons give rise to Co’s chemistry: losing the two 4s electrons (plus one 3d) yields Co^2+ or Co^3+. In periodic trends, cobalt sits between iron (Fe) and nickel (Ni). Its covalent atomic radius is about 125 picometers, between Fe and Ni. The Pauling electronegativity is ~1.88 (lower than Ni’s 1.91, slightly higher than Fe’s 1.83). The first ionization energy of cobalt is about 760 kJ/mol comparable to iron’s (762 kJ/mol) and higher than nickel’s (737 kJ/mol). These values reflect moderate effective nuclear charge: cobalt holds its valence electrons fairly tightly, which correlates with its ability to form Co^2+ and Co^3+. The arrangement of electrons and shell structure means cobalt usually behaves as a typical 3d-transition metal, forming octahedral complexes and showing narrow variations in size and energy across its period.

Isotopes and Nuclear Properties. Natural cobalt consists of a single stable isotope, ^59Co (100% abundance), which has nuclear spin 7/2 and is NMR-active No naturally occurring radioisotopes exist. The most important artificial isotope is ^60Co (half-life 5.27 years) ^60Co decays by β^– emission to stable ^60Ni and emits two high-energy γ-rays (1.17 and 1.33 MeV). These penetrating gamma rays make ^60Co invaluable in medicine and industry for radiotherapy and sterilization. Other radioisotopes include ^57Co (half-life 271.8 days, decays by electron capture to ^57Fe, emitting a 122 keV γ-ray used in diagnostic imaging and Mössbauer spectroscopy) and ^58Co (70.9 days, β^– decay). None of these has a long enough half-life for geologic dating, so cobalt is not used in radiometric age determinations. In summary, ^59Co (stable) dominates natural cobalt, and ^60Co is the notable radioactive source produced in reactors for gamma applications.

Allotropes and Typical Compounds. Elemental cobalt metal has two main allotropes (crystal structures). At ambient temperature it adopts the α-cobalt form, which is hexagonal close-packed (hcp). Above about 700 °C it transforms to β-cobalt with a face-centered cubic (fcc) lattice (similar to the structures of ruthenium and osmium in their respective α/β phases). There are no distinct molecular allotropes like in carbon or sulfur – the allotropes here are simply different metal lattices.

Cobalt forms many compounds, especially with oxidation states +2 and +3. They are often vividly colored. Common types include:

Oxides: Cobalt(II) oxide, CoO, is a greenish solid (basic oxide). Cobalt(II,III) oxide, Co₃O₄, is a black spinel‐structure compound (mixed oxidation Co^2+ and Co^3+) used in catalysts. (Co₂O₃ is rare and decomposes on heating.)
Hydroxides/Sulfides: Cobalt(II) hydroxide Co(OH)₂ is an insoluble pink solid. Cobalt sulfides (CoS, CoS₂) occur in nature (e.g. cobaltite CoAsS), and CoS₂ (cobaltite) is structurally similar to pyrite.
Halides: Cobalt(II) chloride CoCl₂ is a common salt: the anhydrous form is blue, and the hydrated hexahydrate is pink. Cobalt(III) chloride CoCl₃ can be made by oxidation but is unstable. Other halides include CoF₂, CoBr₂, CoI₂ (all +2 salts), and a few Co(III) halides (CoF₃ exists).
Complexes: Cobalt forms numerous coordination compounds. For example, [Co(NH₃)₆]³⁺ (hexamminecobalt(III)) was one of the early Werner complexes. In aqueous solution Co²⁺ is usually hexaaquo Co(H₂O)₆]²⁺, pink); it can be oxidized to Co³⁺ complexes (often green). Cobalt carbonyl Co₂(CO)₈ is a yellow liquid (a metal carbonyl of mixed Co^0/Co^+). Organometallic cobaltocene, Co(C₅H₅)₂ (a ferrocene analog), is a well-known paramagnetic sandwich compound.
Minerals and Pigments: Important cobalt ores include arsenides/sulfides like cobaltite (CoAsS) and skutterudite (CoAs₃). Many cobalt compounds give blue or green pigments: cobalt aluminate (CoAl₂O₄, “cobalt blue”) is a deep blue ceramic pigment, cobalt silicate gives greens, and cobalt glass (smalt) was used historically. These arise from Co²⁺ or Co²⁺/³⁺ chemistry in silicates and phosphates in glazes. Co-containing glazes and glass have been used since antiquity to impart vivid blues and greens.

Physical Properties. Metallic cobalt is dense and refractory. It has a density of about 8.90 g/cm³ at 20 °C. Its melting point is very high: 1768 K (1495 °C), and its boiling point around 3200 K (≈2927 °C). At room temperature α-cobalt is hexagonal close-packed (hcp); the high-temperature β phase is face-centered cubic (fcc). Cobalt is hard (Mohs hardness ~5), tough and moderately ductile when alloyed, though in pure form it is quite brittle. It conducts heat and electricity like a typical metal (thermal conductivity ≈ 100 W·m⁻¹·K⁻¹, electrical resistivity about 6.0×10^–8 Ω·m at 20 °C).

Importantly, cobalt is ferromagnetic at ordinary conditions. Its saturation magnetization is on the order of 1.6 tesla. It retains magnetism up to a Curie temperature of roughly 1130 °C (1403 K) much higher than iron’s Curie point. This high-temperature ferromagnetism contributes to its use in magnets and magnetic recording materials. No strong optical absorption lines of cobalt are generally notable, but atomic emission of cobalt produces sharp blue-green spectral lines (used in analytical spectrometry). In summary, cobalt’s physical profile is that of a strong, high-melt, ferromagnetic metal with the dense, lustrous character of transition metals.

Chemical Reactivity and Trends. Cobalt is moderately reactive chemically. In air at room temperature it does not burn, but slowly forms a thin oxide film. Finely divided cobalt powder can oxidize more readily and even ignite upon heating. Cobalt metal does not react with water, but it dissolves in strong acids: for example, Co + 2HCl → CoCl₂ + H₂. In this way cobalt will liberate hydrogen from acids. Its standard electrode potential (Co²⁺/Co) is about –0.28 V, which is between that of iron and nickel; this means cobalt is less easily oxidized than zinc or iron, but more so than copper. It will corrode in acid, although not as vigorously as very active metals.

Cobalt’s oxides and hydroxides are basic to amphoteric. CoO and Co(OH)₂ dissolve in acids to give Co²⁺. The mixed oxide Co₃O₄ (Co^IICO^III₂O₄) will dissolve in both acids (forming Co²⁺/Co³⁺) and strongly alkaline solutions (forming cobaltate complexes), so it has borderline amphoteric character. Cobalt compounds exhibit the expected acid-base behavior of transition metal salts: cobalt(II) salts in water form slightly acidic solutions (Co²⁺ hydrolyzes weakly), while cobalt(III) oxides/hydroxides are more basic.

In terms of oxidation-reduction, Co(II) is relatively stable but can be further oxidized to Co(III) by strong oxidants (e.g. concentrated HNO₃ or Cl₂). Cobalt(III) compounds (d^6, usually low-spin) are strong oxidizers themselves and typically require complexation (as in vitamin B12 or coordination compounds) to be isolated. For example, cobalt(III) hexammine [Co(NH₃)₆]³⁺ and the tetrachlorocobaltate [CoCl₄]²⁻ demonstrate Co^3+ complexes. Cobalt also forms stable lower-valent organometallic compounds (e.g. Co₂(CO)₈, cobaltocene) under special conditions.

Cobalt does not show dramatically unusual chemistry in the periodic trends: like its neighbors iron and nickel, it forms +2 salts that are often colored (pink or blue for Co²⁺), and +3 complexes that are typically more inert. It does not exhibit strong halogenation to form stable CoF4, CoCl4 (the tetrahydrates of Co²⁺ are known, CoCl₅⁻ etc only under exotic conditions). Overall, cobalt’s reactivity places it in the middle of the transition series: it easily forms Co²⁺, can access Co³⁺ with strong conditions, and its oxides dissolve in acids but do not passivate the metal. Its compounds all trend as expected for a 3d^7 metal ion.

Occurrence and Production. Cobalt is relatively rare in the Earth's crust and is usually mined as a byproduct of other metals. Its cosmic abundance is moderate (~3000 parts per billion by mass in the universe, ~4 parts per million in the Sun) but on Earth it averages only ~30 ppm in crustal rocks Cobalt is most often found combined with iron, nickel, copper, and arsenic in various ores. Common cobalt-bearing minerals include cobaltite (CoAsS), skutterudite (CoAs₃), and erythrite (Co₃(AsO₄)₂·8H₂O). Cobalt is also present in many nickel and copper sulfide ores (for example, the pentlandite ores of nickel often contain a few percent cobalt).

Major global cobalt reserves are in the African Copperbelt (Democratic Republic of Congo and Zambia) and Morocco. The world’s largest producer by far is the DRC (over 60–70% of mined cobalt), followed by countries like Russia, Australia, Canada, Cuba, and Madagascar. Nearly all cobalt is recovered as a byproduct from processing nickel and copper ores. For example, nickel-copper sulfides are roasted or smelted to produce a mixed matte; this matte (or leach solution) is then treated chemically to extract cobalt. Cobalt-rich ores (like cobaltite) can be treated by smelting and hydrometallurgy: roasted to CoO and then leached with acid to yield cobalt salts. The refined cobalt is often produced via precipitation of hydroxides or sulfides, or via solvent extraction systems. An older method (the Mond process) was used historically: cobalt carbonyl Co₂(CO)₈ is formed from impure Co, then decomposed to pure cobalt and CO. Today, most refinery capacity is in Asia (notably China) which imports raw materials for processing.

Cobalt also occurs in minor amounts in the sea (bound to manganese nodules and crusts on the ocean floor, along with Ni and Cu). Meteorites rich in iron-nickel often contain higher cobalt. Overall, cobalt’s supply is sensitive to market and geopolitical factors because it is not mined for itself but follows the production of other metals.

Applications and Technology. Modern industry relies heavily on cobalt’s distinctive properties. Key applications include:

Superalloys: Cobalt-based alloys (often with chromium, nickel, and refractory metals) maintain strength at high temperatures. They are used in jet engine turbine blades, gas turbines, and rocket engines. Cobalt superalloys (e.g. “stellite” and Udimet) resist creep and corrosion at temperatures where other metals would fail.
Permanent Magnets: Cobalt is crucial in high-performance magnets. Alnico alloys (Al–Ni–Co) use about 10–20% Co and were first permanent magnets available (Alnico magnets). More recently, samarium–cobalt (SmCo) rare-earth magnets (with ~30–40% Co) provide strong permanent magnetism that resists demagnetization at high temperature
Lithium-Ion Batteries: Cobalt is a key cathode component in lithium-ion batteries. Early Li-ion cells used lithium cobalt oxide (LiCoO₂) cathodes. Today’s advanced batteries often use layered oxides with nickel, manganese and cobalt (NMC) or nickel, cobalt, aluminium (NCA) to balance energy density and stability. Roughly two-thirds of mined cobalt goes into rechargeable batteries (smartphones, laptops, EVs) because cobalt improves cycle life and safety.
Radioisotope Sources: ^60Co is irradiated cobalt used in medical radiotherapy units (gamma knives) and industrial irradiation. It sterilizes medical equipment and food (killing bacteria), and provides gamma-ray sources for radiography and food irradiation machines Another isotope, ^57Co, is used in medical diagnostics and as a calibration source for gamma detectors.
Pigments and Dyes: Cobalt compounds produce vivid blues and greens. Cobalt aluminate (CoAl₂O₄) is the classic “cobalt blue” pigment in paints and ceramics. Cobalt(II) silicate yields cobalt-green, and other cobalt salts make violet or red dyes. Historically, cobalt pigments have been prized (Egyptian blue glass, Chinese/Tang porcelain glazes, Renaissance stained glass).
Catalysts: Cobalt catalysts are used in chemical processing. Cobalt–molybdenum catalysts (on alumina) remove sulfur from petroleum (hydrodesulfurization). Cobalt metal catalysts (sometimes on carbon) are used in Fischer–Tropsch synthesis converting syngas (CO + H₂) into liquid hydrocarbons. Cobalt compounds also catalyze hydrogenation reactions in fine chemicals.
Hard Metals and Alloys: Tungsten carbide (WC) cutting tools and wear components use cobalt (typically 6–12%) as a metallic binder to hold carbide grains. Cobalt-chromium alloys are used in medical implants (e.g. stents, artificial hips/knees) for their hardness and biocompatibility. Cobalt adds strength and corrosion resistance to steels and superalloys used in drill bits and high-speed tools.
Other Uses: Cobalt salts serve as drying agents (‘driers’) in paints and inks, accelerating polymerization. Cobalt is used in electroplating for a smooth, wear-resistant coating. Its compounds are also used to make magnetic recording media and, in some cases, high-performance glasses and ceramics. Some specialized inorganic cells (like silver-zinc batteries) had trace cobalt.

In all these applications, cobalt’s high melting point, hardness, magnetic stability, and ability to undergo multiple oxidation-reduction cycles are exploited.

Biology, Environment, and Safety. Cobalt is a trace element in biology but is toxic at higher levels. Its only known biological role is as the central metal in vitamin B₁₂ (cobalamin). All animals (including humans) require cobalt only in this organic form; the body needs only a few micrograms of B₁₂ per day. (An average adult stores on the order of 1–2 mg of B₁₂, lasting months to a year without additional intake Cobalt as inorganic metal or salt is not used by the body. Deficiency of vitamin B₁₂ (and thus cobalt) causes pernicious anemia and neurological problems. Some plants (especially in tropical soils) require cobalt for nitrogen-fixing bacteria.

Exposure to cobalt (metal dust or salts) can be harmful. Cobalt dust and fume are respiratory irritants – inhalation can cause asthma-like symptoms and a form of lung disease (“hard metal disease”) Skin contact may cause dermatitis in sensitive individuals. Cobalt compounds (particularly soluble salts like cobalt chloride or sulfate) are moderately toxic by ingestion and can cause heart problems (cardiomyopathy) and thyroid enlargement. Historical episodes (like cobalt in beer additives in the mid-20th century) showed high doses damage the heart. In 2018 the World Health Organization did not set a specific drinking-water limit for cobalt, but water guidelines are on the order of 0.05 mg/L.

Cobalt metal is considered a possible carcinogen when inhaled chronically (IARC classifies cobalt metal with tungsten carbide as Group 2A – probably carcinogenic). Occupational exposure limits (OSHA PEL) are around 0.05–0.1 mg/m³ for dust/fume. In the environment, cobalt binds strongly to soil and organic matter and does not bioaccumulate significantly. Trace cobalt is normal in food and water; excess industrial release can elevate local levels. Radiologically, ^60Co released from nuclear tests once contaminated the environment; today such releases are controlled. Cobalt-60 is also a potent radionuclide hazard if shielding and handling precautions are not followed.

Safety measures for cobalt handling include good ventilation, dust control, and protective clothing. Consumer exposure (e.g. from batteries) is very low. In summary, cobalt is essential in tiny amounts for life (via B₁₂), but its heavy-metal and radioactive forms require careful handling to avoid health risks.

History and Etymology. Cobalt has a long history of use but was recognized as a unique element in the 18th century. Deep blue glass and glazes colored by cobalt minerals have been found in Egyptian artifacts (Tutankhamen’s tomb, 14th century BC) and ancient Chinese pottery However, miners often mistook cobalt ores for bismuth or lead ore, since smelting them produced toxic fumes of arsenic. In the 16th century, German miners near Freiberg, Saxony related the ores to a mischievous “kobold” (goblin) spirit due to their disappointing silver yield

Pure cobalt metal was first isolated by Swedish chemist Georg Brandt around 1735 Brandt showed that “cobalt” was a new element, distinct from bismuth or nickel. He derived the name “cobalt” from the German word Kobold, referencing the goblins of mining folklore. This was the first element chemists identified after the classical metals of antiquity. In the 19th century, advances in chemistry and metallurgy led to systematic study of cobalt compounds and alloys.

In the 20th century, cobalt became strategically important. Cobalt alloys were used in jet engines and military equipment; Cobalt-60 was discovered (in the 1940s) as a potent gamma-ray source. Cobalt blue and other pigments saw huge demand in ceramics and glass, especially in Victorian times. The development of Lithium-ion batteries and superalloys in the late 20th/early 21st century has further highlighted cobalt’s technological role. Today, cobalt’s name reminds us of its mining folklore origin (“kobold” goblin) and its enduring identity as a transition metal essential to both industry and biology.

Property	Cobalt (Co)
Atomic number (Z)	27
Symbol	Co
Atomic weight	58.9332 u (≈58.93)
Group, Period, Block	9, 4, d (transition metal)
Standard state (20 °C)	Solid (silvery-gray metal)
Density (20 °C)	8.90 g/cm³ (8900 kg/m³)
Melting point	1768 K (1495 °C)
Boiling point	3200 K (≈2927 °C)
Crystal structure	hcp (α-Co) at RT; fcc at >700°C
Electrical & thermal	Good conductor; k ≈100 W·m⁻¹·K⁻¹
Electron configuration	[Ar] 3d^7 4s^2
Electronegativity	1.88 (Pauling scale)
First ionization energy	~760 kJ/mol (≈7.88 eV)
Common oxidation states	+2, +3
Stable isotopes	^59Co (100% of natural Co)
Magnetic ordering	Ferromagnetic (Curie ≈1390 K)
Discovery	Georg Brandt, 1735