What actually causes cobalt chloride to change color when its heated?


The water molecules attached to the cobalt II chloride crystals evaporates. What you have after heating is anhydrous cobalt II choride (that is, cobalt II chloride minus water).

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Cobalt(II) chloride Cobaltous chloride
Cobalt dichloride Cl[Co]Cl InChI=1S/2ClH.Co/h2*1H;/q;;+2/p-2Yes 
Key: GVPFVAHMJGGAJG-UHFFFAOYSA-LYes  InChI=1/2ClH.Co/h2*1H;/q;;+2/p-2
140 °C (monohydrate)
100 °C (dihydrate)
86 °C (hexahydrate) 1049 °C Cobalt(II) chloride is an inorganic compound of cobalt and chlorine, with the formula CoCl2. It is usually supplied as the hexahydrate CoCl2·6H2O, which is one of the most commonly used cobalt compounds in the laboratory. The hexahydrate is deep purple in color, whereas the anhydrous form is sky blue. A blend would be mauve. Because of the ease of the hydration/dehydration reaction, and the resulting color change, cobalt chloride is used as an indicator for water in desiccants. Niche uses include its role in organic synthesis and electroplating objects with cobalt metal. Cobalt(II) chloride gives a blue-green color in a flame. Aqueous solutions of both CoCl2 and the hydrate contain the species [Co(H2O)6]2+. They also contain chloride ions. In the solid state CoCl2·6H2O consists of the molecule trans-[CoCl2(H2O)4] and two molecules of water of crystallization. This species dissolves readily in water and alcohol. Concentrated aqueous solutions are red at room temperature but become blue when heated. CoCl2·6H2O is deliquescent and the anhydrous salt CoCl2 is hygroscopic, readily converting to the hydrate. Hydrated cobalt chloride is prepared from cobalt(II) hydroxide or cobalt(II) carbonate and hydrochloric acid: Upon heating, the hexahydrate dehydrates in a stepwise manner. Generally, aqueous solutions of cobalt(II) chlorides behave like other cobalt(II) salts since these solutions consist of the [Co(H2O)6]2+ ion regardless of the anion. Such solutions give a precipitate of CoS upon treatment with S2H. CoCl2·6H2O and CoCl2 are weak Lewis acids that react to give adducts that are usually either octahedral or tetrahedral. With pyridine (), one obtains the octahedral complex: With the bulky ligand triphenylphosphine (), tetrahedral complexes result: The anionic complex CoCl42–: The [CoCl4]2- ion is the blue ion that forms upon addition of hydrochloric acid to aqueous solutions of hydrated cobalt chloride, which are pink. In the laboratory, cobalt(II) chloride serves as a common precursor to other cobalt compounds. Reaction of the anhydrous compound with sodium cyclopentadienide gives cobaltocene. This 19-electron species is a good reducing agent, being readily oxidised to the yellow 18-electron cobaltacenium cation. Reaction of 1-norbonyllithium with the CoCl2·THF in pentane produces the brown, thermally stable cobalt(IV) tetralkyl — a rare example of a stable transition metal/saturated alkane compound, different products are obtained in other solvents. In the presence of ammonia or amines, cobalt(II) is readily oxidised by atmospheric oxygen to give a variety of cobalt(III) complexes. For example, the presence of ammonia triggers the oxidation of cobalt(II) chloride to hexamminecobalt(III) chloride: The reaction is often performed in the presence of charcoal as a catalyst, or hydrogen peroxide is employed in place of air. Other highly basic ligands including carbonate, acetylacetonate, and oxalate induce the formation of Co(III) derivatives. Simple carboxylates and halides do not. Unlike Co(II) complexes, Co(III) complexes are very slow to exchange ligands, so they are said to be kinetically inert. The German chemist Alfred Werner was awarded the Nobel prize in 1913 for his studies on a series of these cobalt(III) compounds, work that led to an understanding of the structures of such coordination compounds. The existence of cobalt(III) chloride, CoCl3, is disputed, although this compound is listed in some compendia. According to Greenwood and Earnshaw, the only stable binary compounds of cobalt and the halogens, excluding CoF3, are the dihalides. Stated differently, CoCl2 is unreactive toward Cl2][. The stability of Co(III) in solution is considerably increased in the presence of ligands of greater Lewis basicity than chloride, such as amines.
Cupric chloride Cl[Cu]Cl [Cu+2].[Cl-].[Cl-] InChI=1S/2ClH.Cu/h2*1H;/q;;+2/p-2Yes 
Key: ORTQZVOHEJQUHG-UHFFFAOYSA-LYes  InChI=1/2ClH.Cu/h2*1H;/q;;+2/p-2/rCl2Cu/c1-3-2
100 °C (dehydration of dihydrate) 993 °C (anhydrous, decomp) Copper(II) chloride is the chemical compound with the formula CuCl2. This is a light brown solid, which slowly absorbs moisture to form a blue-green dihydrate. The copper(II) chlorides are some of the most common copper(II) compounds, after copper sulfate. Anhydrous CuCl2 adopts a distorted cadmium iodide structure. In this motif, the copper centers are octahedral. Most copper(II) compounds exhibit distortions from idealized octahedral geometry due to the Jahn-Teller effect, which in this case describes the localization of one d-electron into a molecular orbital that is strongly antibonding with respect to a pair of chloride ligands. In CuCl2·2H2O, the copper again adopts a highly distorted octahedral geometry, the Cu(II) centers being surrounded by two water ligands and four chloride ligands, which bridge asymmetrically to other Cu centers. Copper(II) chloride is paramagnetic. Of historical interest, CuCl2·2H2O was used in the first electron paramagnetic resonance measurements by Yevgeny Zavoisky in 1944. Aqueous solution prepared from copper(II) chloride contain a range of copper(II) complexes depending on concentration, temperature, and the presence of additional chloride ions. These species include blue color of [Cu(H2O)6]2+ and yellow or red color of the halide complexes of the formula [CuCl2+x]x−. Copper(II) hydroxide precipitates upon treating copper(II) chloride solutions with base: Partial hydrolysis gives copper oxychloride, Cu2Cl(OH)3, a popular fungicide. CuCl2 is a mild oxidant. It decomposes to CuCl and 2Cl at 1000 °C: CuCl2 reacts with several metals to produce copper metal or copper(I) chloride with oxidation of the other metal. To convert copper(II) chloride to copper(I) derivatives, it can be convenient to reduce an aqueous solution with sulfur dioxide as the reductant: CuCl2 reacts with HCl or other chloride sources to form complex ions: the red CuCl3−, and the yellow CuCl42−. Some of these complexes can be crystallized from aqueous solution, and they adopt a wide variety of structures. Copper(II) chloride also forms a variety of coordination complexes with ligands such as pyridine and triphenylphosphine oxide: However "soft" ligands such as phosphines (e.g., triphenylphosphine), iodide, and cyanide as well as some tertiary amines induce reduction to give copper(I) complexes. Copper(II) chloride is prepared commercially by the action of chlorination of copper: It can also be generated by treatment of the hydroxide, oxide, or copper(II) carbonate with hydrochloric acid. Electrolysis of aqueous sodium chloride with copper electrodes produces (among other things) a blue-green foam that can be collected and converted to the hydrate. Copper(II) chloride can also be prepared by mixing dilute hydrochloric acid with copper(II) sulfate. Anhydrous CuCl2 may be prepared directly by union of the elements, copper and chlorine. CuCl2 may be purified by crystallization from hot dilute hydrochloric acid, by cooling in a 2CaCl-ice bath. Copper(II) chloride occurs naturally as the very rare mineral tolbachite and the dihydrate . Both are found near fumaroles. More common are mixed oxyhydroxide-chlorides like atacamite Cu2(OH)3Cl, arising among Cu ore beds oxidation zones in arid climate (also known from some altered slags). A major industrial application for copper(II) chloride is as a co-catalyst with palladium(II) chloride in the Wacker process. In this process, ethene (ethylene) is converted to ethanal (acetaldehyde) using water and air. During the reaction, 2PdCl is reduced to Pd, and the CuCl2 serves to re-oxidize this back to PdCl2. Air can then oxidize the resultant CuCl back to CuCl2, completing the cycle. The overall process is: Copper(II) chloride is used as a catalyst in a variety of processes that produce chlorine by oxychlorination. The Deacon process takes place at about 400 to 450 °C in the presence of a copper chloride: Copper(II) chloride catalyzes the chlorination in the production of vinyl chloride and dichloroethane. Copper(II) chloride is used in the Copper–chlorine cycle in which it splits steam into a copper oxygen compound and hydrogen chloride, and is later recovered in the cycle from the electrolysis of copper(I) chloride. Copper(II) chloride has some highly specialized applications in the synthesis of organic compounds. It effects chlorination of aromatic hydrocarbons- this is often performed in the presence of aluminium oxide. It is able to chlorinate the alpha position of carbonyl compounds: This reaction is performed in a polar solvent such as dimethylformamide (DMF), often in the presence of lithium chloride, which accelerates the reaction. CuCl2, in the presence of oxygen, can also oxidize phenols. The major product can be directed to give either a quinone or a coupled product from oxidative dimerization. The latter process provides a high-yield route to 1,1-binaphthol: Such compounds are intermediates in the synthesis of BINAP and its derivatives Copper(II) chloride dihydrate promotes the hydrolysis of acetonides, i.e., for deprotection to regenerate diols or aminoalcohols, as in this example (where TBDPS = -butyldiphenylsilyltert): CuCl2 also catalyses the free radical addition of sulfonyl chlorides to alkenes; the alpha-chlorosulfone may then undergo elimination with base to give a vinyl sulfone product.][ Copper(II) chloride is also used in pyrotechnics as a blue/green coloring agent. In a flame test, copper chlorides, like all copper compounds, emit green-blue. In humidity indicator cards (HICs), cobalt-free brown to azure (copper(II) chloride base) HICs can be found on the market. In 1998, the European Community (EC) classified items containing cobalt(II) chloride of 0.01 to 1% w/w as T (Toxic), with the corresponding R phrase of R49 (may cause cancer if inhaled). As a consequence, new cobalt-free humidity indicator cards have been developed that contain copper. It is toxic and only concentrations below 5 ppm are allowed in drinking water by the US Environmental Protection Agency.
Cobalt(II) hydroxide Cobaltous hydroxide, cobalt hydroxide, cobaltous hydrate [Co+2].[OH-].[OH-] InChI=1S/Co.2H2O/h;2*1H2/q+2;;/p-2Yes 
Key: ASKVAEGIVYSGNY-UHFFFAOYSA-LYes  InChI=1/Co.2H2O/h;2*1H2/q+2;;/p-2
Key: ASKVAEGIVYSGNY-NUQVWONBAS 168 °C (decomp) Cobalt(II) hydroxide or cobaltous hydroxide is the chemical compound composed of cobalt and the hydroxide ion with the formula Co(OH)2. It occurs in two forms, either as a rose-red powder, which is the more stable of the two forms, or as bluish-green powder. It has the or cadmium iodide crystal structure. It finds use as a drying agent for paints, varnishes and inks, in the preparation of other cobalt compounds, as a catalyst and in the manufacture of battery electrodes. Cobalt(II) hydroxide is precipitated when an alkaline hydroxide is added to an aqueous solution of Co2+ ions: Cobalt(II) hydroxide decomposes to cobalt(II) oxide at 168 °C under vacuum and is oxidized by air to form cobalt(III) hydroxide, Co(OH)3. The thermal decomposition product in air above 300 °C is 4O3Co. Like iron(II) hydroxide, cobalt(II) hydroxide is primarily a basic hydroxide, although it does form the weakly acidic reddish hexaaquacobalt(II) ion, [Co(H2O)6]2+, in acidic aqueous solutions. In strong bases, cobalt(II) hydroxide accepts additional hydroxide ions to form dark blue cobaltates(II) [Co(OH)4]2- and [Co(OH)6]4-.
[Co](Br)Br InChI=1S/2BrH.Co/h2*1H;/q;;+2/p-2Yes 
Key: BZRRQSJJPUGBAA-UHFFFAOYSA-LYes  InChI=1/2BrH.Co/h2*1H;/q;;+2/p-2
47 °C (hexahydrate) Cobalt(II) bromide (CoBr2) is an inorganic compound. It is a red solid that is soluble in water, used primarily as a catalyst in some processes. When anhydrous, cobalt(II) bromide appears as green crystals. The hexahydrate loses four waters of crystallization molecules at 100 °C forming the dihydrate: Further heating to 130 °C produces the anhydrous form: The anhydrous form melts at 678 °C. At higher temperatures, cobalt(II) bromide reacts with oxygen, forming cobalt(II,III) oxide and bromine vapor. Cobalt(II) bromide can be prepared as a hydrate by the reaction of cobalt hydroxide with hydrobromic acid: Anhydrous cobalt(II) bromide may be prepared through the direct reaction of elemental cobalt and liquid bromine. The classical coordination compound bromopentaamminecobalt(III) bromide is prepared by oxidation of a solution of cobalt(II) bromide in aqueous ammonia. Triphenylphosphine complexes of cobalt(II) bromide have been used as a catalysts in organic synthesis. Exposure to large amounts of cobalt(II) can cause cobalt poisoning. Bromide is also mildly toxic.
Cobalt(II) sulfate [Co+2].[O-]S([O-])(=O)=O InChI=1S/Co.H2O4S/c;1-5(2,3)4/h;(H2,1,2,3,4)/q+2;/p-2Yes 
Key: KTVIXTQDYHMGHF-UHFFFAOYSA-LYes  InChI=1/Co.H2O4S/c;1-5(2,3)4/h;(H2,1,2,3,4)/q+2;/p-2
96.8 °C (heptahydrate) 420 °C (heptahydrate) Cobalt(II) sulfate is the inorganic compound with the formula CoSO4. It is the divalent cobalt salt of sulfuric acid. The most common form of cobalt sulfate are the hydrates CoSO4.7H2O and CoSO4.H2O. Cobalt(II) sulfate and its hydrates are some of the most commonly available salts of cobalt. Cobalt(II) sulfate appears as red monoclinic crystals that melt around 100 °C and become anhydrous at 250 °C. It is soluble in water, slightly soluble in ethanol, and especially soluble in methanol. It forms by the reaction of metallic cobalt, its oxide, hydroxide, or carbonate with sulfuric acid. Cobalt is obtained from ores via the sulfate in some cases. Cobalt(II) sulfate is used in the preparation of pigments, as well as in the manufacture of other cobalt salts. Cobalt pigment is used in porcelains and glass. Cobalt(II) sulfate is used in storage batteries and electroplating baths, sympathetic inks, and as an additive to soils and animal feeds. For these purposes, the cobalt sulfate is produced by treating cobalt oxide with sulfuric acid. Cobalt is essential for most higher forms of life, but more than a few milligrams each day is harmful. Rarely have poisonings resulted from cobalt compounds. Upon inhalation of salts, there is some evidence for carcinogenicity.
1195 °C Cobalt sulfide is the name for chemical compounds with a formula CoxSy. Well-characterized species include minerals with the formula CoS2 and Co3S4, and the synthetic material Co9S8. In combination with molybdenum, the sulfides of cobalt are used as catalysts for the industrial process called hydrodesulfurization, which is implemented on a large scale in refineries. Cobalt sulfides precipitate when aqueous solutions of cobalt(II) ions are treated with hydrogen sulfide. This reaction is useful in the purification of cobalt from its ores as well as in qualitative inorganic analysis. In general, the sulfides of cobalt are black, semiconducting, insoluble in water, and nonstoichiometric. They react with strong acids to release hydrogen sulfide gas again. They are weak reducing agents and can be oxidized to cobalt sulfate][. Cobalt sulfide exists in two forms: alpha and beta. The best defined sulfides of cobalt occur as minerals. The rare mineral cattierite has the stoichiometry CoS2. It is isostructural with iron pyrite, featuring disulfide groups, i.e. Co2+S22-. Linnaeite, also rare, has the formula Co3S4 and crystallizes in the spinel motif.
A humidity indicator is a moisture-sensitive chemical that changes color when the indicated relative humidity is exceeded. Some chemicals are:

Metal halides are compounds between metals and halogens. Some, such as sodium chloride are ionic, while others are covalently bonded. Covalently bonded metal halides may be discrete molecules, such as uranium hexafluoride, or they may form polymeric structures, such as palladium chloride.

Sodium chloride crystal structure

Cobalt is a chemical element with symbol Co and atomic number 27. Like nickel, cobalt in the Earth's crust is found only in chemically combined form, save for small deposits found in alloys of natural meteoric iron. The free element, produced by reductive smelting, is a hard, lustrous, silver-gray metal.

Cobalt-based blue pigments (cobalt blue) have been used since ancient times for jewelry and paints, and to impart a distinctive blue tint to glass, but the color was later thought by alchemists to be due to the known metal bismuth. Miners had long used the name kobold ore (German for goblin ore) for some of the blue-pigment producing minerals; they were named because they were poor in known metals, and gave poisonous arsenic-containing fumes upon smelting. In 1735, such ores were found to be reducible to a new metal (the first discovered since ancient times), and this was ultimately named for the kobold.

Hydrate is a term used in inorganic chemistry and organic chemistry to indicate that a substance contains water. The chemical state of the water varies widely between hydrates, some of which were so labeled before their chemical structure was understood.

In organic chemistry, a hydrate is a compound formed by the addition of water or its elements to another molecule. For example, ethanol, CH3–CH2–OH, can be considered as a hydrate of ethene, CH2=CH2, formed by the addition of H to one C and OH to the other C. A molecule of water may be eliminated, for example by the action of sulfuric acid. Another example is chloral hydrate, CCl3–CH(OH)2, which can be formed by reaction of water with chloral, CCl3–CH=O.

Desiccation is the state of extreme dryness, or the process of extreme drying. A desiccant is a hygroscopic (attracts and holds water) substance that induces or sustains such a state in its local vicinity in a moderately sealed container.

A desiccator is a heavy glass or plastic container used in practical chemistry for making or keeping small amounts of materials very dry. The material is placed on a shelf, and a drying agent or desiccant, such as dry silica gel or anhydrous sodium hydroxide, is placed below the shelf.


Chemistry, a branch of physical science, is the study of the composition, properties and behavior of matter. Chemistry is chiefly concerned with atoms and their interactions with other atoms - for example, the properties of the chemical bonds formed between atoms to create chemical compounds. As well as this, interactions including atoms and other phenomena - electrons and various forms of energy - are considered, such as photochemical reactions, oxidation-reduction reactions, changes in phases of matter, and separation of mixtures. Finally, properties of matter such as alloys or polymers are considered.

Chemistry is sometimes called "the central science" because it bridges other natural sciences like physics, geology and biology with each other. Chemistry is a branch of physical science but distinct from physics.



Dietary minerals (also known as mineral nutrients) are the chemical elements required by living organisms, other than the four elements carbon, hydrogen, nitrogen, and oxygen present in common organic molecules. The term is archaic, as it describes chemical elements rather than actual minerals.

Minerals in order of abundance in the human body include the seven major minerals calcium, phosphorus, potassium, sulfur, sodium, chlorine, and magnesium. Important "trace" or minor minerals, necessary for mammalian life, include iron, cobalt, copper, zinc, molybdenum, iodine, and selenium (see below for detailed discussion).


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