Electronegativity, symbol χ, is a chemical property that describes the tendency of an atom to attract a shared pair of electrons towards itself. An atom's electronegativity is affected by both its atomic number and the distance at which its valence electrons reside from the charged nucleus; the higher the associated electronegativity number, the more an atom or a substituent group attracts electrons towards itself. On the most basic level, electronegativity is determined by factors like the nuclear charge and the number/location of other electrons present in the atomic shells; the opposite of electronegativity is electropositivity: a measure of an element's ability to donate electrons. The term "electronegativity" was introduced by Jöns Jacob Berzelius in 1811, though the concept was known before that and was studied by many chemists including Avogadro. In spite of its long history, an accurate scale of electronegativity was not developed until 1932, when Linus Pauling proposed an electronegativity scale, which depends on bond energies, as a development of valence bond theory.
It has been shown to correlate with a number of other chemical properties. Electronegativity cannot be directly measured and must be calculated from other atomic or molecular properties. Several methods of calculation have been proposed, although there may be small differences in the numerical values of the electronegativity, all methods show the same periodic trends between elements; the most used method of calculation is that proposed by Linus Pauling. This gives a dimensionless quantity referred to as the Pauling scale, on a relative scale running from around 0.7 to 3.98. When other methods of calculation are used, it is conventional to quote the results on a scale that covers the same range of numerical values: this is known as an electronegativity in Pauling units; as it is calculated, electronegativity is not a property of an atom alone, but rather a property of an atom in a molecule. Properties of a free atom include ionization electron affinity, it is to be expected that the electronegativity of an element will vary with its chemical environment, but it is considered to be a transferable property, to say that similar values will be valid in a variety of situations.
Caesium is the least electronegative element in the periodic table, while fluorine is most electronegative. Francium and caesium were both assigned 0.7. However, francium's ionization energy is known to be higher than caesium's, in accordance with the relativistic stabilization of the 7s orbital, this in turn implies that francium is in fact more electronegative than caesium. Pauling first proposed the concept of electronegativity in 1932 as an explanation of the fact that the covalent bond between two different atoms is stronger than would be expected by taking the average of the strengths of the A–A and B–B bonds. According to valence bond theory, of which Pauling was a notable proponent, this "additional stabilization" of the heteronuclear bond is due to the contribution of ionic canonical forms to the bonding; the difference in electronegativity between atoms A and B is given by: | χ A − χ B | = − 1 / 2 E d − E d + E d 2 where the dissociation energies, Ed, of the A–B, A–A and B–B bonds are expressed in electronvolts, the factor −1⁄2 being included to ensure a dimensionless result.
Hence, the difference in Pauling electronegativity between hydrogen and bromine is 0.73 As only differences in electronegativity are defined, it is necessary to choose an arbitrary reference point in order to construct a scale. Hydrogen was chosen as the reference, as it forms covalent bonds with a large variety of elements: its electronegativity was fixed first at 2.1 revised to 2.20. It is necessary to decide which of the two elements is the more electronegative; this is done using "chemical intuition": in the above example, hydrogen bromide dissolves in water to form H+ and Br− ions, so it may be assumed that bromine is more electronegative than hydrogen. However, in principle, since the same electronegativities should be obtained for any two bonding compounds, the data are in fact overdetermined, the signs are unique once a reference point is fixed. To calculate Pauling electronegativity for an element, it
Acid dissociation constant
An acid dissociation constant, Ka, is a quantitative measure of the strength of an acid in solution. It is the equilibrium constant for a chemical reaction known as dissociation in the context of acid–base reactions. K a =; the chemical species HA, A−, H+ are said to be in equilibrium when their concentrations do not change with the passing of time, because both forward and backward reactions are occurring at the same fast rate. The chemical equation for acid dissociation can be written symbolically as: HA ↽ − − ⇀ A − + H + where HA is a generic acid that dissociates into A−, the conjugate base of the acid and a hydrogen ion, H+, it is implicit in this definition that the quotient of activity coefficients, Γ, Γ = γ A − γ H + γ A H is a constant that can be ignored in a given set of experimental conditions. For many practical purposes it is more convenient to discuss the logarithmic constant, pKa p K a = − log 10 The more positive the value of pKa, the smaller the extent of dissociation at any given pH —that is, the weaker the acid.
A weak acid has a pKa value in the approximate range −2 to 12 in water. For a buffer solution consisting of a weak acid and its conjugate base, pKa can be expressed as: p K a = pH − log 10 The pKa for a weak monoprotic acid is conveniently determined by potentiometric titration with a strong base to the equivalence point and taking the pH value measured at one-half this volume as being equal to pKa; that is because at this half equivalence point, the number of moles of strong base added is one-half the number of moles of weak acid present, while the concentrations of the conjugate base and the remaining weak acid are the same. Acids with a pKa value of less than about −2 are said to be strong acids. In water, the dissociation of a strong acid in dilute solutions is complete such that the final concentration of the undissociated acid final is low. Consider a strong monoprotic acid, such as HCl; because of their 1:1 ratio, the final concentration of the conjugate base, final, is taken to be equal to the concentration of the hydronium ion, which can be directly measured by a pH meter.
For strong monoprotic acids like HCl, final and are both nearly equal to the initial concentration of initial placed into solution. With conventional acid-base titration methods it is difficult to measure the pH of a strong acid solution and, hence, to determine the or final, with a sufficient number of significant figures to and compute the low values encountered for final, which can be as low as 10-9 mol per liter for some strong acids. Furthermore, if 100% dissociation is assumed, final is zero and the fraction within parenthesis in the equation above becomes undefined; because the second expression on the right-hand side of the above equation is therefore indeterminable by conventional titration methods, the entire equation is not as useful a means of experimentally measuring pKa for strong acids as it is for weak acids. However, pKa and/or Ka values for strong acids can be estimated by theoretical means, such as computing gas phase dissociation constants and using Gibbs free energies of solvation for the molecular anions.
It is possible to use spectroscopy in some cases to determine the ratio of the concentrations of the conjugate base produced and the undissociated acid. For example, the Raman spectra of dilute nitric acid solutions contain signals of the nitrate ion and as the solutions become more concentrated signals of undissociated nitric acid molecules emerge; the acid dissociation constant for an acid is a direct consequence of the underlying thermodynamics of the dissociation reaction. The value of the pKa changes with temperature and can be understood qualitatively based on Le Châtelier's principle: when the reaction is endothermic, Ka increases and pKa decreases with
A carboxylic acid is an organic compound that contains a carboxyl group. The general formula of a carboxylic acid is R–COOH, with R referring to the rest of the molecule. Carboxylic acids occur widely. Important examples include acetic acid. Deprotonation of a carboxyl group gives a carboxylate anion. Important carboxylate salts are soaps. Carboxylic acids are identified by their trivial names, they have the suffix -ic acid. IUPAC-recommended names exist. For example, butyric acid is butanoic acid by IUPAC guidelines. For nomenclature of complex molecules containing a carboxylic acid, the carboxyl can be considered position one of the parent chain if there are other substituents, for example, 3-chloropropanoic acid. Alternately, it can be named as a "carboxy" or "carboxylic acid" substituent on another parent structure, for example, 2-carboxyfuran; the carboxylate anion of a carboxylic acid is named with the suffix -ate, in keeping with the general pattern of -ic acid and -ate for a conjugate acid and its conjugate base, respectively.
For example, the conjugate base of acetic acid is acetate. Carboxylic acids are polar; because they are both hydrogen-bond acceptors and hydrogen-bond donors, they participate in hydrogen bonding. Together the hydroxyl and carbonyl group forms the functional group carboxyl. Carboxylic acids exist as dimers in nonpolar media due to their tendency to "self-associate". Smaller carboxylic acids are soluble in water, whereas higher carboxylic acids have limited solubility due to the increasing hydrophobic nature of the alkyl chain; these longer chain acids tend to be rather soluble in less-polar solvents such as ethers and alcohols. Hydrophobic carboxylic acids react aqueous sodium hydroxide to give water soluble sodium salts. For example, enathic acid has a small solubility in water, but its sodium salt is soluble in water: Carboxylic acids tend to have higher boiling points than water, not only because of their increased surface area, but because of their tendency to form stabilised dimers through hydrogen bonds.
For boiling to occur, either the dimer bonds must be broken or the entire dimer arrangement must be vaporised, both of which increase the enthalpy of vaporization requirements significantly. Carboxylic acids are Brønsted -- Lowry acids, they are the most common type of organic acid. Carboxylic acids are weak acids, meaning that they only dissociate into H3O+ cations and RCOO− anions in neutral aqueous solution. For example, at room temperature, in a 1-molar solution of acetic acid, only 0.4% of the acid are dissociated. Electron-withdrawing substituents, such as -CF3 group, give stronger acids. Electron-donating substituents give weaker acids Deprotonation of carboxylic acids gives carboxylate anions; each of the carbon–oxygen bonds in the carboxylate anion has a partial double-bond character. The carbonyl carbon's partial positive charge is weakened by the -1/2 negative charges on the 2 oxygen atoms. Carboxylic acids have strong sour odors. Esters of carboxylic acids tend to have pleasant odors, many are used in perfume.
Carboxylic acids are identified as such by infrared spectroscopy. They exhibit a sharp band associated with vibration of the C–O vibration bond between 1680 and 1725 cm−1. A characteristic νO–H band appears as a broad peak in the 2500 to 3000 cm−1 region. By 1H NMR spectrometry, the hydroxyl hydrogen appears in the 10–13 ppm region, although it is either broadened or not observed owing to exchange with traces of water. Many carboxylic acids are produced industrially on a large scale, they are pervasive in nature. Esters of fatty acids are the main components of lipids and polyamides of aminocarboxylic acids are the main components of proteins. Carboxylic acids are used in the production of polymers, pharmaceuticals and food additives. Industrially important carboxylic acids include acetic acid and methacrylic acids, adipic acid, citric acid, ethylenediaminetetraacetic acid, fatty acids, maleic acid, propionic acid, terephthalic acid. In general, industrial routes to carboxylic acids differ from those used on smaller scale because they require specialized equipment.
Carbonylation of alcohols as illustrated by the Cativa process for production of acetic acid. Formic acid is prepared by a different carbonylation pathway starting from methanol. Oxidation of aldehydes with air using cobalt and manganese catalysts; the required aldehydes are obtained from alkenes by hydroformylation. Oxidation of hydrocarbons using air. For simple alkanes, this method is inexpensive but not selective enough to be useful. Allylic and benzylic compounds undergo more selective oxidations. Alkyl groups on a benzene ring are oxidized to the carboxylic acid, regardless of its chain length. Benzoic acid from toluene, terephthalic acid from para-xylene, phthalic acid from ortho-xylene are illustrative large-scale conversions. Acrylic acid is generated from propene. Base-cata
Citric acid is a weak organic acid that has the chemical formula C6H8O7. It occurs in citrus fruits. In biochemistry, it is an intermediate in the citric acid cycle, which occurs in the metabolism of all aerobic organisms. More than a million tons of citric acid are manufactured every year, it is used as an acidifier, as a flavoring and chelating agent. A citrate is a derivative of citric acid. An example of the former, a salt is trisodium citrate; when part of a salt, the formula of the citrate ion is written as C6H5O3−7 or C3H5O3−3. Citric acid exists in greater than trace amounts in a variety of fruits and vegetables, most notably citrus fruits. Lemons and limes have high concentrations of the acid; the concentrations of citric acid in citrus fruits range from 0.005 mol/L for oranges and grapefruits to 0.30 mol/L in lemons and limes. Within species, these values vary depending on the cultivar and the circumstances in which the fruit was grown. Industrial-scale citric acid production first began in 1890 based on the Italian citrus fruit industry, where the juice was treated with hydrated lime to precipitate calcium citrate, isolated and converted back to the acid using diluted sulfuric acid.
In 1893, C. Wehmer discovered. However, microbial production of citric acid did not become industrially important until World War I disrupted Italian citrus exports. In 1917, American food chemist James Currie discovered certain strains of the mold Aspergillus niger could be efficient citric acid producers, the pharmaceutical company Pfizer began industrial-level production using this technique two years followed by Citrique Belge in 1929. In this production technique, still the major industrial route to citric acid used today, cultures of A. niger are fed on a sucrose or glucose-containing medium to produce citric acid. The source of sugar is corn steep liquor, hydrolyzed corn starch or other inexpensive sugary solutions. After the mold is filtered out of the resulting solution, citric acid is isolated by precipitating it with calcium hydroxide to yield calcium citrate salt, from which citric acid is regenerated by treatment with sulfuric acid, as in the direct extraction from citrus fruit juice.
In 1977, a patent was granted to Lever Brothers for the chemical synthesis of citric acid starting either from aconitic or isocitrate/alloisocitrate calcium salts under high pressure conditions. This produced citric acid in near quantitative conversion under what appeared to be a reverse non-enzymatic Krebs cycle reaction. In 2007, worldwide annual production stood at 1,600,000 tons. More than 50% of this volume was produced in China. More than 50% was used as an acidity regulator in beverages, some 20% in other food applications, 20% for detergent applications and 10% for related applications other than food, such as cosmetics, pharmaceutics and in the chemical industry. Citric acid was first isolated in 1784 by the chemist Carl Wilhelm Scheele, who crystallized it from lemon juice, it can exist either as a monohydrate. The anhydrous form crystallizes from hot water, while the monohydrate forms when citric acid is crystallized from cold water; the monohydrate can be converted to the anhydrous form at about 78 °C.
Citric acid dissolves in absolute ethanol at 15 °C. It decomposes with loss of carbon dioxide above about 175 °C. Citric acid is considered to be a tribasic acid, with pKa values, extrapolated to zero ionic strength, of 5.21, 4.28 and 2.92 at 25 °C. The pKa of the hydroxyl group has been found, by means of 13C NMR spectroscopy, to be 14.4. The speciation diagram shows that solutions of citric acid are buffer solutions between about pH 2 and pH 8. In biological systems around pH 7, the two species present are the citrate ion and mono-hydrogen citrate ion; the SSC 20X hybridization buffer is an example in common use. Tables compiled for biochemical studies are available. On the other hand, the pH of a 1 mM solution of citric acid will be about 3.2. The pH of fruit juices from citrus fruits like oranges and lemons depends on the citric acid concentration, being lower for higher acid concentration and conversely. Acid salts of citric acid can be prepared by careful adjustment of the pH before crystallizing the compound.
See, for example, sodium citrate. The citrate ion forms complexes with metallic cations; the stability constants for the formation of these complexes are quite large because of the chelate effect. It forms complexes with alkali metal cations. However, when a chelate complex is formed using all three carboxylate groups, the chelate rings have 7 and 8 members, which are less stable thermodynamically than smaller chelate rings. In consequence, the hydroxyl group can be deprotonated, forming part of a more stable 5-membered ring, as in ammonium ferric citrate, 5Fe2·2H2O. Citric acid can be esterified at one or more of the carboxylic acid functional groups on the molecule, to form any of a variety of mono-, di-, tri-, mixed esters. Citrate is an intermediate in the TCA cycle, a central metabolic pathway for animals and bacteria. Citrate synthase catalyzes the condensation of oxaloacetate with acetyl CoA to form citrate. Citrate acts as the substrate for aconitase and is converted into aconitic acid.
The cycle ends with regeneration of oxaloacetate. This series
The SN2 reaction is a type of reaction mechanism, common in organic chemistry. In this mechanism, one bond is broken and one bond is formed synchronously, i.e. in one step. SN2 is a kind of nucleophilic substitution reaction mechanism. Since two reacting species are involved in the slow step, this leads to the term substitution nucleophilic or SN2, the other major kind is SN1. Many other more specialized mechanisms describe substitution reactions; the reaction type is so common that it has other names, e.g. "bimolecular nucleophilic substitution", or, among inorganic chemists, "associative substitution" or "interchange mechanism". The reaction most occurs at an aliphatic sp3 carbon center with an electronegative, stable leaving group attached to it, a halide atom; the breaking of the C–X bond and the formation of the new bond occur through a transition state in which a carbon under nucleophilic attack is pentacoordinate, sp2 hybridised. The nucleophile attacks the carbon at 180° to the leaving group, since this provides the best overlap between the nucleophile's lone pair and the C–X σ* antibonding orbital.
The leaving group is pushed off the opposite side and the product is formed with inversion of the tetrahedral geometry at the central atom. If the substrate under nucleophilic attack is chiral this leads to inversion of configuration, called a Walden inversion. In an example of the SN2 reaction, the attack of Br− on an ethyl chloride results in ethyl bromide, with chloride ejected as the leaving group.: SN2 attack occurs if the backside route of attack is not sterically hindered by substituents on the substrate. Therefore, this mechanism occurs at unhindered primary and secondary carbon centres. If there is steric crowding on the substrate near the leaving group, such as at a tertiary carbon centre, the substitution will involve an SN1 rather than an SN2 mechanism. Four factors affect the rate of the reaction: The substrate plays the most important part in determining the rate of the reaction; this is because the nucleophile attacks from the back of the substrate, thus breaking the carbon-leaving group bond and forming the carbon-nucleophile bond.
Therefore, to maximise the rate of the SN2 reaction, the back of the substrate must be as unhindered as possible. Overall, this means that methyl and primary substrates react the fastest, followed by secondary substrates. Tertiary substrates do not participate in SN2 reactions, because of steric hindrance. Structures that can form stable cations by simple loss of the leaving group, for example, as a resonance-stabilized carbocation, are likely to react via an SN1 pathway in competition with SN2. Like the substrate, steric hindrance affects the nucleophile's strength; the methoxide anion, for example, is both a strong base and nucleophile because it is a methyl nucleophile, is thus much unhindered. Tert-Butoxide, on the other hand, is a strong base, but a poor nucleophile, because of its three methyl groups hindering its approach to the carbon. Nucleophile strength is affected by charge and electronegativity: nucleophilicity increases with increasing negative charge and decreasing electronegativity.
For example, OH− is a better nucleophile than water, I− is a better nucleophile than Br−. In a polar aprotic solvent, nucleophilicity increases up a column of the periodic table as there is no hydrogen bonding between the solvent and nucleophile. I − would therefore be a weaker nucleophile than Br −. Verdict - A strong/anionic nucleophile always favours SN2 manner of nucleophillic substitution; the solvent affects the rate of reaction because solvents may or may not surround a nucleophile, thus hindering or not hindering its approach to the carbon atom. Polar aprotic solvents, like tetrahydrofuran, are better solvents for this reaction than polar protic solvents because polar protic solvents will hydrogen bond to the nucleophile, hindering it from attacking the carbon with the leaving group. A polar aprotic solvent with low dielectric constant or a hindered dipole end will favour SN2 manner of nucleophilic substitution reaction. Examples: DMSO, DMF, acetone etc. In polar aprotic solvent, nucleophilicity parallels basicity.
The stability of the leaving group as an anion and the strength of its bond to the carbon atom both affect the rate of reaction. The more stable the conjugate base of the leaving group is, the more that it will take the two electrons of its bond to carbon during the reaction. Therefore, the weaker the leaving group is as a conjugate base, thus the stronger its corresponding acid, the better the leaving group. Examples of good leaving groups are therefore the halides and tosylate, whereas HO− and H2N− are not; the rate of an SN2 reaction is second order, as the rate-determining step depends on the nucleophile concentration, as well as the concentration of substrate. R = kThis is a key difference between the SN2 mechanisms. In the SN1 reaction the nucleophile attacks after the rate-limiting step is over, whereas in SN2 the nucleophile forces off the leaving group in the limiting step. In other words, the rate of SN1 reactions depend only on the concentration of the substrate while the SN2 reaction rate depends on the concentration of both the substrate and nucleophile.
It has been shown that except in uncommon primary and secondary
Formate is the anion derived from formic acid. Its formula is represented in various equivalent ways: HCOO− or CHOO− or HCO2−, it is the product of deprotonation of formic acid. It is the simplest carboxylate anion. A formate is a ester of formic acid. Formate is reversibly oxidized by the enzyme formate dehydrogenase from Desulfovibrio gigas: HCO2− → CO2 + H+ + 2 e− Formate esters have the formula ROCH. Many form spontaneously; the most important formate ester is methyl formate, produced as an intermediate en route to formic acid. Methanol and carbon monoxide react in the presence of a strong base, such as sodium methoxide: CH3OH + CO → HCO2CH3Hydrolysis of methyl formate gives formic acid and regenerates methanol: HCO2CH3 → HCO2H + CH3OHFormic acid is used for many applications in industry. Formate esters are fragrant or have distinctive odors. Compared to the more common ethyl esters, formate esters are less used commercially because they are less stable. Ethyl formate is found in some confectionaries.
Formate salts have the formula Mx. Such salts are prone to decarboxylation. For example, hydrated nickel formate decarboxylates at about 200 °C to give finely powdered nickel metal: Ni22 → Ni + 2 CO2 + 2 H2O + H2Such fine powders are useful as hydrogenation catalysts. Ethyl formate, CH3CH2 sodium formate, Na potassium formate, K caesium formate, Cs.
An alkoxide is the conjugate base of an alcohol and therefore consists of an organic group bonded to a negatively charged oxygen atom. They are written as RO −. Alkoxides are strong bases and, when R is not good nucleophiles and good ligands. Alkoxides, although not stable in protic solvents such as water, occur as intermediates in various reactions, including the Williamson ether synthesis. Transition metal alkoxides are used for coatings and as catalysts. Enolates are unsaturated alkoxides derived by deprotonation of a C-H bond adjacent to a ketone or aldehyde; the nucleophilic center for simple alkoxides is located on the oxygen, whereas the nucleophilic site on enolates is delocalized onto both carbon and oxygen sites. Phenoxides are close relatives of the alkoxides, in which the alkyl group is replaced by a derivative of benzene. Phenol is more acidic than a typical alcohol, they are, however easier to handle, yield derivatives that are more crystalline than those of the alkoxides. Alkali metal alkoxides are oligomeric or polymeric compounds when the R group is small.
The alkoxide anion is a good bridging ligand, thus many alkoxides feature M2O or M3O linkages. In solution, the alkali metal derivatives exhibit strong ion-pairing, as expected for the alkali metal derivative of a basic anion. Alkoxides can be produced by several routes starting from an alcohol. Reducing metals react directly with alcohols to give the corresponding metal alkoxide; the alcohol serves as an acid, hydrogen is produced as a by-product. A classic case is sodium methoxide produced by the addition of sodium metal to methanol: 2 CH3OH + 2 Na → 2 CH3ONa + H2Other alkali metals can be used in place of sodium, most alcohols can be used in place of methanol. Another similar reaction occurs when an alcohol is reacted with a metal hydride such as NaH; the metal hydride removes the hydrogen atom from the hydroxyl group and forms a negatively charged alkoxide ion. Titanium tetrachloride reacts with alcohols to give the corresponding tetraalkoxides, concomitant with the evolution of hydrogen chloride: TiCl4 + 4 2CHOH → Ti4 + 4 HClThe reaction can be accelerated by the addition of a base, such as a tertiary amine.
Many other metal and main group halides can be used instead of titanium, for example SiCl4, ZrCl4, PCl3. Many alkoxides are prepared by salt-forming reactions from a metal chloride and sodium alkoxide: n NaOR + MCln → Mn + n NaClSuch reactions are favored by the lattice energy of the NaCl, purification of the product alkoxide is simplified by the fact that NaCl is insoluble in common organic solvents. Many alkoxides can be prepared by anodic dissolution of the corresponding metals in water-free alcohols in the presence of electroconductive additive; the metals may be etc.. The conductive additive may be quaternary ammonium halide, or other; some examples of metal alkoxides obtained by this technique: Ti4, Nb210, Ta210, 2, Re2O36, Re4O612, Re4O610. The alkoxide ion can react with a primary alkyl halide in an SN2 reaction to form a methyl ether. Metal alkoxides hydrolyse with water according to the following equation: 2 LnMOR + H2O → 2O + 2 ROHwhere R is an organic substituent and L is an unspecified ligand A well-studied case is the irreversible hydrolysis of titanium ethoxide: 1/n n + 2 H2O → TiO2 + 4 HOCH2CH3By controlling the stoichiometry and steric properties of the alkoxide, such reactions can be arrested leading to metal-oxy-alkoxides, which are oligonuclear complexes.
Other alcohols can be employed in place of water. In this way one alkoxide can be converted to another, the process is properly referred to as alcoholysis; the position of the equilibrium can be controlled by the acidity of the alcohol. More the alcoholysis can be controlled by selectively evaporating the more volatile component. In this way, ethoxides can be converted to butoxides. In the transesterification process, metal alkoxides react with esters to bring about an exchange of alkyl groups between metal alkoxide and ester. With the metal alkoxide complex in focus, the result is the same as for alcoholysis, namely the replacement of alkoxide ligands, but at the same time the alkyl groups of the ester are changed, which can be the primary goal of the reaction. Sodium methoxide, for example, is used for this purpose, a reaction, relevant to the production of "bio-diesel". Many metal alkoxide compounds feature oxo-ligands. Oxo-ligands arise via the hydrolysis accidentally, via ether elimination: 2 LnMOR → 2O + R2OAdditionally, low valent metal alkoxides are susceptible to oxidation by air Characteristically, transition metal alkoxides are polynuclear, they contain more than one metal.
Alkoxides are sterically undemanding and basic ligands that tend to bridge metals. Upon the isomorphic substitution of metal atoms close in properties crystalline complexes of variable composition are formed; the metal ratio in such compounds can vary over a broad range. For instance, the substitution of molybdenum and tungsten for rhenium in the complexes Re4O6−y12+y allo