In chemistry, molecular oxohalides (oxyhalides) are a group of chemical compounds in which both oxygen and halogen atoms are attached to another chemical element A in a single molecule. They have the general formula AOmXn, where X is a halogen. Known oxohalides have fluorine (F), chlorine (Cl), bromine (Br), and/or iodine (I) in their molecules. The element A may be a main group element, a transition element, a rare earth element or an actinide. The term oxohalide, or oxyhalide, may also refer to minerals and other crystalline substances with the same overall chemical formula, but having an ionic structure.
Oxohalides can be seen as compounds intermediate between oxides and halides. There are three general methods of synthesis: [1]
In addition, various oxohalides can be made by halogen exchange reactions and this reaction can also lead to the formation of mixed oxohalides such as POFCl2 and CrO2FCl.
In relation to the oxide or halide, for a given oxidation state of an element A, if two halogen atoms replace one oxygen atom, or vice versa, the overall charge on the molecule is unchanged and the coordination number of the central atom decreases by one. For example, both phosphorus oxychloride (POCl3) and phosphorus pentachloride, (PCl5) are neutral covalent compounds of phosphorus in the +5 oxidation state. If an oxygen atom is simply replaced by a halogen atom the charge increases by +1, but the coordination number is unchanged. This is illustrated by the reaction of a mixture of a chromate or dichromate salt and potassium chloride with concentrated sulfuric acid.
The chromyl chloride produced has no electrical charge and is a volatile covalent molecule that can be distilled out of the reaction mixture. [2]
Oxohalides of elements in high oxidation states are strong oxidizing agents, with oxidizing power similar to the corresponding oxide or halide. Most oxohalides are easily hydrolyzed. For example, chromyl chloride is hydrolyzed to chromate in the reverse of the synthetic reaction, above. The driving force for this reaction is the formation of A-O bonds which are stronger than A-Cl bonds. This gives a favourable enthalpy contribution to the Gibbs free energy change for the reaction [3]
Many oxohalides can act as Lewis acids. This is particularly so with oxohalides of coordination number 3 or 4 which, in accepting one or more electron pairs from a Lewis base, become 5- or 6-coordinate. Oxohalide anions such as [VOCl4]2− can be seen as acid-base complexes of the oxohalide (VOCl2) with more halide ions acting as Lewis bases. Another example is VOCl2 which forms the trigonal bipyramidal complex VOCl2(N(CH3)3)2 with the base trimethylamine. [4]
The vibrational spectra of many oxohalides have been assigned in detail. They give useful information on relative bond strengths. For example, in CrO2F2, the Cr–O stretching vibrations are at 1006 cm−1 and 1016 cm−1 and the Cr–F stretching vibrations are at 727 cm−1 and 789 cm−1. The difference is much too large to be due to the different masses of O and F atoms. Rather, it shows that the Cr–O bond is much stronger than the Cr–F bond. M–O bonds are generally considered to be double bonds and this is backed up by measurements of M–O bond lengths. It implies that the elements A and O are chemically bound together by a σ bond and a π bond. [5]
Oxohalides of elements in high oxidation states are intensely coloured owing to ligand to metal charge transfer (LMCT) transitions. [6]
Carbon forms oxohalides COX2, X = F, Br, and the very toxic phosgene (X = Cl), which is produced industrially by a carbon-catalyzed reaction of carbon monoxide with chlorine. It is a useful reagent in organic chemistry for the formation of carbonyl compounds. [7] For example:
Silicon tetrafluoride reacts with water to yield poorly-characterized oxyfluoride polymers, but slow and careful reaction at -196 °C yields the oxyfluoride hexafluorodisiloxane as well. [8]
Nitrogen forms two series of oxohalides with nitrogen in oxidation states 3, NOX, X = F, Cl, Br and 5, NO2X, X = F, Cl. They are made by halogenation of nitrogen oxides. Note that NO2F is isoelectronic with the nitrate ion, NO−3. Only oxohalides of phosphorus(V) are known. [9]
Sulfur forms oxohalides [10] in oxidation state +4, such as thionyl chloride, SOCl2 and oxidation state +6, such as sulfuryl fluoride (SO2F2), sulfuryl chloride (SO2Cl2), and thionyl tetrafluoride (SOF4). All are easily hydrolyzed. Indeed, thionyl chloride can be used as a dehydration agent as the water molecules are converted into gaseous products, leaving behind the anhydrous solid chloride. [11]
Selenium and tellurium form similar compounds and also the oxo-bridged species F5AOAF5 (A = S, Se, Te). They are non-linear with the A-O-A angle of 142.5, 142.4 and 145.5° for S, Se and Te, respectively. [12] The tellurium anion F5TeO−, known as teflate, is a large and rather stable anion, useful for forming stable salts with large cations. [11]
The halogens form various oxofluorides with formulae XO2F (chloryl fluoride), XO3F (perchloryl fluoride) and XOF3 with X = Cl, Br and I. IO2F3 and IOF5 are also known. [13]
Xenon forms xenon oxytetrafluoride (XeOF4), xenon dioxydifluoride (XeO2F2) and xenon oxydifluoride (XeOF2).
A selection of known oxohalides of transition metals is shown below, and more detailed lists are available in the literature. [15] X indicates various halides, most often F and Cl.
Oxidation state | oxohalides |
---|---|
3 | VOCl, VOBr, [16] FeOCl |
4 | [TiOCl4]2−, Cl3TiOTiCl3, VOCl2 , [VOCl4]2− |
5 | VOX3, VO2F , [CrOF4]−, [CrOF5]2−, MnOCl3, TcOCl3, VOF3 , VOCl3 , NbOCl3 |
6 | CrO2Cl2 , [CrO3Cl]−, CrOF4 , ReOX4, ReO2F2, OsOF4, CrO2F2 , MoOF4 , MoOCl4 , MoO2Cl2 , MoO2F2 , WO2Cl2 , WO2F2, WOF4 , WOCl4 |
7 | MnO3F , ReOF5, ReO2F3 , ReO3F , ReO3Cl , OsOF5 |
8 | OsO2F4, OsO3F2 |
High oxidation states of the metal are dictated by the fact that oxygen is a strong oxidizing agent, as is fluorine. Bromine and iodine are relatively weak oxidizing agents, so fewer oxobromides and oxoiodides are known. Structures for compounds with d0 configuration are predicted by VSEPR theory. Thus, CrO2Cl2 is tetrahedral, OsO3F2 is trigonal bipyramidal, XeOF4 is square pyramidal and OsOF5 is octahedral. [18] The d1 complex ReOCl4 is square pyramidal.
The compounds [Ta2OX10]2− and [M2OCl10]4− (M = W, Ru, Os) have two MX5 groups joined by a bridging oxygen atom. [19] Each metal has an octahedral environment. The unusual linear M−O−M structure can be rationalized in terms of molecular orbital theory, indicating the presence of dπ — pπ bonding between the metal and oxygen atoms. [20] Oxygen bridges are present in more complex configurations like M(cp)2(OTeF5)2 (M = Ti, Zr, Hf, Mo or W; cp = cyclopentadienyl, η5-C5H5) [21] or [AgOTeF5-(C6H5CH3)2]2. [17]
In the actinide series, uranyl compounds such as uranyl chloride (UO2Cl2) and [UO2Cl4]2− are well known and contain the linear UO2 moiety. Similar species exist for neptunium and plutonium.
Bismuth oxochloride (BiOCl, bismoclite) is a rare example of a mineral oxohalide. The crystal structure has a tetragonal symmetry and can be thought of as consisting of layers of Cl−, Bi3+ and O2− ions, in the order Cl-Bi-O-Bi-Cl-Cl-Bi-O-Bi-Cl. This layered, graphite-like structure results in a relatively low hardness of bismoclite (Mohs 2–2.5) and most other oxohalide minerals. [22] Those other minerals include terlinguaite Hg2OCl, formed by the weathering of mercury-containing minerals. [23] Mendipite, Pb3O2Cl2, formed from an original deposit of lead sulfide in a number of stages is another example of a secondary oxohalide mineral.
The elements iron, antimony, bismuth and lanthanum form oxochlorides of general formula MOCl. MOBr and MOI are also known for Sb and Bi. Many of their crystal structures have been determined. [24]
Bromine is a chemical element; it has symbol Br and atomic number 35. It is a volatile red-brown liquid at room temperature that evaporates readily to form a similarly coloured vapour. Its properties are intermediate between those of chlorine and iodine. Isolated independently by two chemists, Carl Jacob Löwig and Antoine Jérôme Balard, its name was derived from the Ancient Greek βρῶμος (bromos) meaning "stench", referring to its sharp and pungent smell.
Chlorine is a chemical element; it has symbol Cl and atomic number 17. The second-lightest of the halogens, it appears between fluorine and bromine in the periodic table and its properties are mostly intermediate between them. Chlorine is a yellow-green gas at room temperature. It is an extremely reactive element and a strong oxidising agent: among the elements, it has the highest electron affinity and the third-highest electronegativity on the revised Pauling scale, behind only oxygen and fluorine.
The halogens are a group in the periodic table consisting of six chemically related elements: fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and the radioactive elements astatine (At) and tennessine (Ts), though some authors would exclude tennessine as its chemistry is unknown and is theoretically expected to be more like that of gallium. In the modern IUPAC nomenclature, this group is known as group 17.
In chemistry, the oxidation state, or oxidation number, is the hypothetical charge of an atom if all of its bonds to other atoms were fully ionic. It describes the degree of oxidation of an atom in a chemical compound. Conceptually, the oxidation state may be positive, negative or zero. While fully ionic bonds are not found in nature, many bonds exhibit strong ionicity, making oxidation state a useful predictor of charge.
The octet rule is a chemical rule of thumb that reflects the theory that main-group elements tend to bond in such a way that each atom has eight electrons in its valence shell, giving it the same electronic configuration as a noble gas. The rule is especially applicable to carbon, nitrogen, oxygen, and the halogens; although more generally the rule is applicable for the s-block and p-block of the periodic table. Other rules exist for other elements, such as the duplet rule for hydrogen and helium, and the 18-electron rule for transition metals.
In chemistry, an interhalogen compound is a molecule which contains two or more different halogen atoms and no atoms of elements from any other group.
The chemical element nitrogen is one of the most abundant elements in the universe and can form many compounds. It can take several oxidation states; but the most common oxidation states are -3 and +3. Nitrogen can form nitride and nitrate ions. It also forms a part of nitric acid and nitrate salts. Nitrogen compounds also have an important role in organic chemistry, as nitrogen is part of proteins, amino acids and adenosine triphosphate.
Technetium compounds are chemical compounds containing the chemical element technetium. Technetium can form multiple oxidation states, but often forms in the +4 and +7 oxidation states. Because technetium is radioactive, technetium compounds are extremely rare on Earth.
Bromine compounds are compounds containing the element bromine (Br). These compounds usually form the -1, +1, +3 and +5 oxidation states. Bromine is intermediate in reactivity between chlorine and iodine, and is one of the most reactive elements. Bond energies to bromine tend to be lower than those to chlorine but higher than those to iodine, and bromine is a weaker oxidising agent than chlorine but a stronger one than iodine. This can be seen from the standard electrode potentials of the X2/X− couples (F, +2.866 V; Cl, +1.395 V; Br, +1.087 V; I, +0.615 V; At, approximately +0.3 V). Bromination often leads to higher oxidation states than iodination but lower or equal oxidation states to chlorination. Bromine tends to react with compounds including M–M, M–H, or M–C bonds to form M–Br bonds.
Chromium compounds are compounds containing the element chromium (Cr). Chromium is a member of group 6 of the transition metals. The +3 and +6 states occur most commonly within chromium compounds, followed by +2; charges of +1, +4 and +5 for chromium are rare, but do nevertheless occasionally exist.
Iodine compounds are compounds containing the element iodine. Iodine can form compounds using multiple oxidation states. Iodine is quite reactive, but it is much less reactive than the other halogens. For example, while chlorine gas will halogenate carbon monoxide, nitric oxide, and sulfur dioxide, iodine will not do so. Furthermore, iodination of metals tends to result in lower oxidation states than chlorination or bromination; for example, rhenium metal reacts with chlorine to form rhenium hexachloride, but with bromine it forms only rhenium pentabromide and iodine can achieve only rhenium tetraiodide. By the same token, however, since iodine has the lowest ionisation energy among the halogens and is the most easily oxidised of them, it has a more significant cationic chemistry and its higher oxidation states are rather more stable than those of bromine and chlorine, for example in iodine heptafluoride.
Iron shows the characteristic chemical properties of the transition metals, namely the ability to form variable oxidation states differing by steps of one and a very large coordination and organometallic chemistry: indeed, it was the discovery of an iron compound, ferrocene, that revolutionalized the latter field in the 1950s. Iron is sometimes considered as a prototype for the entire block of transition metals, due to its abundance and the immense role it has played in the technological progress of humanity. Its 26 electrons are arranged in the configuration [Ar]3d64s2, of which the 3d and 4s electrons are relatively close in energy, and thus it can lose a variable number of electrons and there is no clear point where further ionization becomes unprofitable.
There are three sets of Indium halides, the trihalides, the monohalides, and several intermediate halides. In the monohalides the oxidation state of indium is +1 and their proper names are indium(I) fluoride, indium(I) chloride, indium(I) bromide and indium(I) iodide.
There are three sets of gallium halides, the trihalides where gallium has oxidation state +3, the intermediate halides containing gallium in oxidation states +1, +2 and +3 and some unstable monohalides, where gallium has oxidation state +1.
Zinc compounds are chemical compounds containing the element zinc which is a member of the group 12 of the periodic table. The oxidation state of zinc in most compounds is the group oxidation state of +2. Zinc may be classified as a post-transition main group element with zinc(II). Zinc compounds are noteworthy for their nondescript appearance and behavior: they are generally colorless, do not readily engage in redox reactions, and generally adopt symmetrical structures.
Polyhalogen ions are a group of polyatomic cations and anions containing halogens only. The ions can be classified into two classes, isopolyhalogen ions which contain one type of halogen only, and heteropolyhalogen ions with more than one type of halogen.
Fluorine forms a great variety of chemical compounds, within which it always adopts an oxidation state of −1. With other atoms, fluorine forms either polar covalent bonds or ionic bonds. Most frequently, covalent bonds involving fluorine atoms are single bonds, although at least two examples of a higher order bond exist. Fluoride may act as a bridging ligand between two metals in some complex molecules. Molecules containing fluorine may also exhibit hydrogen bonding. Fluorine's chemistry includes inorganic compounds formed with hydrogen, metals, nonmetals, and even noble gases; as well as a diverse set of organic compounds. For many elements the highest known oxidation state can be achieved in a fluoride. For some elements this is achieved exclusively in a fluoride, for others exclusively in an oxide; and for still others the highest oxidation states of oxides and fluorides are always equal.
Many compounds of thorium are known: this is because thorium and uranium are the most stable and accessible actinides and are the only actinides that can be studied safely and legally in bulk in a normal laboratory. As such, they have the best-known chemistry of the actinides, along with that of plutonium, as the self-heating and radiation from them is not enough to cause radiolysis of chemical bonds as it is for the other actinides. While the later actinides from americium onwards are predominantly trivalent and behave more similarly to the corresponding lanthanides, as one would expect from periodic trends, the early actinides up to plutonium have relativistically destabilised and hence delocalised 5f and 6d electrons that participate in chemistry in a similar way to the early transition metals of group 3 through 8: thus, all their valence electrons can participate in chemical reactions, although this is not common for neptunium and plutonium.
Carbon oxohalides are a group of chemical compounds that contain only carbon, oxygen and halogen atoms: fluorine, chlorine, bromine and iodine. They include carbonyl halides, COX2, and oxalyl halides, C2X2O2, where X = F, Cl, Br or I. The halogen atoms X do not have to be identical; they differ in mixed oxohalides. Most combinations of halogens exist but carbonyl iodide, COI2, is unknown. The carbon–oxygen bond length in carbonyl halides (1.13–1.17 Å) is shorter than in other carbonyl compounds such as aldehydes and ketones, carboxylic acids, esters and amides. They are reactive reagents for halogenation, acylation and dehydration reactions.
Gallium compounds are compounds containing the element gallium. These compounds are found primarily in the +3 oxidation state. The +1 oxidation state is also found in some compounds, although it is less common than it is for gallium's heavier congeners indium and thallium. For example, the very stable GaCl2 contains both gallium(I) and gallium(III) and can be formulated as GaIGaIIICl4; in contrast, the monochloride is unstable above 0 °C, disproportionating into elemental gallium and gallium(III) chloride. Compounds containing Ga–Ga bonds are true gallium(II) compounds, such as GaS (which can be formulated as Ga24+(S2−)2) and the dioxan complex Ga2Cl4(C4H8O2)2. There are also compounds of gallium with negative oxidation states, ranging from -5 to -1, most of these compounds being magnesium gallides (MgxGay).
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