A molecule is an electrically neutral group of two or more atoms held together by chemical bonds. Molecules are distinguished from ions by their lack of electrical charge. However, in quantum physics, organic chemistry, biochemistry, the term molecule is used less also being applied to polyatomic ions. In the kinetic theory of gases, the term molecule is used for any gaseous particle regardless of its composition. According to this definition, noble gas atoms are considered molecules as they are monatomic molecules. A molecule may be homonuclear, that is, it consists of atoms of one chemical element, as with oxygen. Atoms and complexes connected by non-covalent interactions, such as hydrogen bonds or ionic bonds, are not considered single molecules. Molecules as components of matter are common in organic substances, they make up most of the oceans and atmosphere. However, the majority of familiar solid substances on Earth, including most of the minerals that make up the crust and core of the Earth, contain many chemical bonds, but are not made of identifiable molecules.
No typical molecule can be defined for ionic crystals and covalent crystals, although these are composed of repeating unit cells that extend either in a plane or three-dimensionally. The theme of repeated unit-cellular-structure holds for most condensed phases with metallic bonding, which means that solid metals are not made of molecules. In glasses, atoms may be held together by chemical bonds with no presence of any definable molecule, nor any of the regularity of repeating units that characterizes crystals; the science of molecules is called molecular chemistry or molecular physics, depending on whether the focus is on chemistry or physics. Molecular chemistry deals with the laws governing the interaction between molecules that results in the formation and breakage of chemical bonds, while molecular physics deals with the laws governing their structure and properties. In practice, this distinction is vague. In molecular sciences, a molecule consists of a stable system composed of two or more atoms.
Polyatomic ions may sometimes be usefully thought of as electrically charged molecules. The term unstable molecule is used for reactive species, i.e. short-lived assemblies of electrons and nuclei, such as radicals, molecular ions, Rydberg molecules, transition states, van der Waals complexes, or systems of colliding atoms as in Bose–Einstein condensate. According to Merriam-Webster and the Online Etymology Dictionary, the word "molecule" derives from the Latin "moles" or small unit of mass. Molecule – "extremely minute particle", from French molécule, from New Latin molecula, diminutive of Latin moles "mass, barrier". A vague meaning at first; the definition of the molecule has evolved. Earlier definitions were less precise, defining molecules as the smallest particles of pure chemical substances that still retain their composition and chemical properties; this definition breaks down since many substances in ordinary experience, such as rocks and metals, are composed of large crystalline networks of chemically bonded atoms or ions, but are not made of discrete molecules.
Molecules are held together by ionic bonding. Several types of non-metal elements exist only as molecules in the environment. For example, hydrogen only exists as hydrogen molecule. A molecule of a compound is made out of two or more elements. A covalent bond is a chemical bond; these electron pairs are termed shared pairs or bonding pairs, the stable balance of attractive and repulsive forces between atoms, when they share electrons, is termed covalent bonding. Ionic bonding is a type of chemical bond that involves the electrostatic attraction between oppositely charged ions, is the primary interaction occurring in ionic compounds; the ions are atoms that have lost one or more electrons and atoms that have gained one or more electrons. This transfer of electrons is termed electrovalence in contrast to covalence. In the simplest case, the cation is a metal atom and the anion is a nonmetal atom, but these ions can be of a more complicated nature, e.g. molecular ions like NH4+ or SO42−. An ionic bond is the transfer of electrons from a metal to a non-metal for both atoms to obtain a full valence shell.
Most molecules are far too small to be seen with the naked eye. DNA, a macromolecule, can reach macroscopic sizes, as can molecules of many polymers. Molecules used as building blocks for organic synthesis have a dimension of a few angstroms to several dozen Å, or around one billionth of a meter. Single molecules cannot be observed by light, but small molecules and the outlines of individual atoms may be traced in some circumstances by use of an atomic force microscope; some of the largest molecules are supermolecules. The smallest molecule is the diatomic hydrogen, with a bond length of 0.74 Å. Effective molecular radius is the size; the table of permselectivity for different substances contains examples. The chemical formula for a molecule uses one line of chemical element symbols and sometimes al
Nucleophile is a chemical species that donates an electron pair to form a chemical bond in relation to a reaction. All molecules or ions with a free pair of electrons or at least one pi bond can act as nucleophiles; because nucleophiles donate electrons, they are by definition Lewis bases. Nucleophilic describes the affinity of a nucleophile to the nuclei. Nucleophilicity, sometimes referred to as nucleophile strength, refers to a substance's nucleophilic character and is used to compare the affinity of atoms. Neutral nucleophilic reactions with solvents such as alcohols and water are named solvolysis. Nucleophiles may take part in nucleophilic substitution, whereby a nucleophile becomes attracted to a full or partial positive charge; the terms nucleophile and electrophile were introduced by Christopher Kelk Ingold in 1933, replacing the terms anionoid and cationoid proposed earlier by A. J. Lapworth in 1925; the word nucleophile is derived from philos for love. In general, in a row across the periodic table, the more basic the ion the more reactive it is as a nucleophile.
Within a series of nucleophiles with the same attacking element, the order of nucleophilicity will follow basicity. Sulfur is in general a better nucleophile than oxygen. Many schemes attempting to quantify relative nucleophilic strength have been devised; the following empirical data have been obtained by measuring reaction rates for a large number of reactions involving a large number of nucleophiles and electrophiles. Nucleophiles displaying the so-called alpha effect are omitted in this type of treatment; the first such attempt is found in the Swain–Scott equation derived in 1953: log 10 = s n This free-energy relationship relates the pseudo first order reaction rate constant, k, of a reaction, normalized to the reaction rate, k0, of a standard reaction with water as the nucleophile, to a nucleophilic constant n for a given nucleophile and a substrate constant s that depends on the sensitivity of a substrate to nucleophilic attack. This treatment results in the following values for typical nucleophilic anions: acetate 2.7, chloride 3.0, azide 4.0, hydroxide 4.2, aniline 4.5, iodide 5.0, thiosulfate 6.4.
Typical substrate constants are 0.66 for ethyl tosylate, 0.77 for β-propiolactone, 1.00 for 2,3-epoxypropanol, 0.87 for benzyl chloride, 1.43 for benzoyl chloride. The equation predicts that, in a nucleophilic displacement on benzyl chloride, the azide anion reacts 3000 times faster than water; the Ritchie equation, derived in 1972, is another free-energy relationship: log 10 = N + where N+ is the nucleophile dependent parameter and k0 the reaction rate constant for water. In this equation, a substrate-dependent parameter like s in the Swain–Scott equation is absent; the equation states that two nucleophiles react with the same relative reactivity regardless of the nature of the electrophile, in violation of the reactivity–selectivity principle. For this reason, this equation is called the constant selectivity relationship. In the original publication the data were obtained by reactions of selected nucleophiles with selected electrophilic carbocations such as tropylium or diazonium cations: or ions based on Malachite green.
Many other reaction types have since been described. Typical Ritchie N+ values are: 0.5 for methanol, 5.9 for the cyanide anion, 7.5 for the methoxide anion, 8.5 for the azide anion, 10.7 for the thiophenol anion. The values for the relative cation reactivities are −0.4 for the malachite green cation, +2.6 for the benzenediazonium cation, +4.5 for the tropylium cation. In the Mayr-Patz equation: log = s The second order reaction rate constant k at 20 °C for a reaction is related to a nucleophilicity parameter N, an electrophilicity parameter E, a nucleophile-dependent slope parameter s; the constant s is defined as 1 with 2-methyl-1-pentene as the nucleophile. Many of the constants have been derived from reaction of so-called benzhydrylium ions as the electrophiles: and a diverse collection of π-nucleophiles:. Typical E values are +6.2 for R = chlorine, +5.90 for R = hydrogen, 0 for R = methoxy and -7.02 for R = dimethylamine. Typical N values with s in parenthesis are -4.47 for electrophilic aromatic substitution to toluene, -0.41 for electrophilic addition to 1-phenyl-2-propene, 0.96 for addition to 2-methyl-1-pentene, -0.13 for reaction with triphenylallylsilane, 3.61 for reaction with 2-methylfuran, +7.48 for reaction with isobutenyltributylstannane and +13.36 for reaction with the enamine 7.
The range of organic reactions include SN2 reactions: With E = -9.15 for the S-methyldibenzothiophenium ion, typical nucleophile values N are 15.63 for piperidine, 10.49 for methoxide, 5.20 for water. In short, nucleophilicities towards sp2 or sp3 centers follow the same pattern. In an effort to unify the above described equations the Mayr equation is rewritten as: log = s E s N ( N + E
In chemistry, bases are substances that, in aqueous solution, release hydroxide ions, are slippery to the touch, can taste bitter if an alkali, change the color of indicators, react with acids to form salts, promote certain chemical reactions, accept protons from any proton donor or contain or displaceable OH− ions. Examples of bases are the hydroxides of the alkaline earth metals; these particular substances produce hydroxide ions in aqueous solutions, are thus classified as Arrhenius bases. For a substance to be classified as an Arrhenius base, it must produce hydroxide ions in an aqueous solution. Arrhenius believed; this makes the Arrhenius model limited, as it cannot explain the basic properties of aqueous solutions of ammonia or its organic derivatives. There are bases that do not contain a hydroxide ion but react with water, resulting in an increase in the concentration of the hydroxide ion. An example of this is the reaction between water to produce ammonium and hydroxide. In this reaction ammonia is the base.
Ammonia and other bases similar to it have the ability to form a bond with a proton due to the unshared pair of electrons that they possess. In the more general Brønsted–Lowry acid–base theory, a base is a substance that can accept hydrogen cations —otherwise known as protons. In the Lewis model, a base is an electron pair donor. In water, by altering the autoionization equilibrium, bases yield solutions in which the hydrogen ion activity is lower than it is in pure water, i.e. the water has a pH higher than 7.0 at standard conditions. A soluble base is called an alkali if it releases OH − ions quantitatively. However, it is important to realize. Metal oxides and alkoxides are basic, conjugate bases of weak acids are weak bases. Bases can be thought of as the chemical opposite of acids. However, some strong acids are able to act as bases. Bases and acids are seen as opposites because the effect of an acid is to increase the hydronium concentration in water, whereas bases reduce this concentration.
A reaction between an acid and a base is called neutralization. In a neutralization reaction, an aqueous solution of a base reacts with an aqueous solution of an acid to produce a solution of water and salt in which the salt separates into its component ions. If the aqueous solution is saturated with a given salt solute, any additional such salt precipitates out of the solution; the notion of a base as a concept in chemistry was first introduced by the French chemist Guillaume François Rouelle in 1754. He noted that acids, which at that time were volatile liquids, turned into solid salts only when combined with specific substances. Rouelle considered that such a substance serves as a "base" for the salt, giving the salt a "concrete or solid form". General properties of bases include: Concentrated or strong bases are caustic on organic matter and react violently with acidic substances. Aqueous solutions or molten bases dissociate in ions and conduct electricity. Reactions with indicators: bases turn red litmus paper blue, phenolphthalein pink, keep bromothymol blue in its natural colour of blue, turn methyl orange yellow.
The pH of a basic solution at standard conditions is greater than seven. Bases are bitter in taste; the following reaction represents the general reaction between a base and water to produce a conjugate acid and a conjugate base: B + H2O ⇌ BH+ + OH−The equilibrium constant, Kb, for this reaction can be found using the following general equation: Kb = /In this equation, both the base and the strong base compete with one another for the proton. As a result, bases that react with water have small equilibrium constant values; the base is weaker. Bases react with acids to neutralize each other at a fast rate both in alcohol; when dissolved in water, the strong base sodium hydroxide ionizes into hydroxide and sodium ions: NaOH → Na+ + OH−and in water the acid hydrogen chloride forms hydronium and chloride ions: HCl + H2O → H3O+ + Cl−When the two solutions are mixed, the H3O+ and OH− ions combine to form water molecules: H3O+ + OH− → 2 H2OIf equal quantities of NaOH and HCl are dissolved, the base and the acid neutralize leaving only NaCl table salt, in solution.
Weak bases, such as baking soda or egg white, should be used to neutralize any acid spills. Neutralizing acid spills with strong bases, such as sodium hydroxide or potassium hydroxide, can cause a violent exothermic reaction, the base itself can cause just as much damage as the original acid spill. Bases are compounds that can neutralize an amount of acids. Both sodium carbonate and ammonia are bases, although neither of these substances contains OH− groups. Both compounds accept H+ when dissolved in protic solvents such as water: Na2CO3 + H2O → 2 Na+ + HCO3− + OH− NH3 + H2O → NH4+ + OH−From this, a pH, or acidity, can be calculated for aqueous solutions of bases. Bases directly act as electron-pair donors themselves: CO32− + H+ → HCO3− NH3 + H+ → NH4+A base is defined as a molecule that has the ability to accept an electron pair bond by entering another atom's valence shell through its possession of one electron pair. There are a limited number of elements that have atoms with the ability to provide a molecule with basic properties
In chemistry, a radical is an atom, molecule, or ion that has an unpaired valence electron. With some exceptions, these unpaired electrons make radicals chemically reactive. Many radicals spontaneously dimerize. Most organic radicals have short lifetimes. A notable example of a radical is the hydroxyl radical, a molecule that has one unpaired electron on the oxygen atom. Two other examples are triplet triplet carbene which have two unpaired electrons. Radicals may be generated in a number of ways. Ionizing radiation, electrical discharges, electrolysis are known to produce radicals. Radicals are intermediates in many chemical reactions, more so than is apparent from the balanced equations. Radicals are important in combustion, atmospheric chemistry, plasma chemistry and many other chemical processes. A large fraction of natural products is generated by radical-generating enzymes. In living organisms, the radicals superoxide and nitric oxide and their reaction products regulate many processes, such as control of vascular tone and thus blood pressure.
They play a key role in the intermediary metabolism of various biological compounds. Such radicals can be messengers in a process dubbed redox signaling. A radical may be otherwise bound. In chemical equations, radicals are denoted by a dot placed to the right of the atomic symbol or molecular formula as follows: C l 2 → U V 2 C l ⋅ Radical reaction mechanisms use single-headed arrows to depict the movement of single electrons: The homolytic cleavage of the breaking bond is drawn with a'fish-hook' arrow to distinguish from the usual movement of two electrons depicted by a standard curly arrow; the second electron of the breaking bond moves to pair up with the attacking radical electron. Radicals take part in radical addition and radical substitution as reactive intermediates. Chain reactions involving radicals can be divided into three distinct processes; these are initiation and termination. Initiation reactions are those, they may involve the formation of radicals from stable species as in Reaction 1 above or they may involve reactions of radicals with stable species to form more radicals.
Propagation reactions are those reactions involving radicals in which the total number of radicals remains the same. Termination reactions are those reactions resulting in a net decrease in the number of radicals. Two radicals combine to form a more stable species, for example: 2Cl·→ Cl2 Radicals can form by breaking of covalent bonds by homolysis; the homolytic bond dissociation energies abbreviated as "ΔH °" are a measure of bond strength. Splitting H2 into 2H•, for example, requires a ΔH ° of +435 kJ·mol-1, while splitting Cl2 into 2Cl• requires a ΔH ° of +243 kJ·mol-1. For weak bonds, homolysis can be induced thermally. Strong bonds require high energy photons or flames to induce homolysis. Radicals or charged species add to non-radicals to give new radicals; this process is the basis of the radical chain reaction. Being prevalent and a diradical, O2 reacts with many organic compounds to generate radicals together with the hydroperoxide radical; this process is related to rancidification of unsaturated fats.
Radicals may be formed by single-electron oxidation or reduction of an atom or molecule. These redox reactions occur in electrochemical cells and in ionization chambers of mass spectrometers. Although radicals are short-lived due to their reactivity, there are long-lived radicals; these are categorized as follows: The prime example of a stable radical is molecular dioxygen. Another common example is nitric oxide. Organic radicals can be long lived if they occur in a conjugated π system, such as the radical derived from α-tocopherol. There are hundreds of examples of thiazyl radicals, which show low reactivity and remarkable thermodynamic stability with only a limited extent of π resonance stabilization. Persistent radical compounds are those whose longevity is due to steric crowding around the radical center, which makes it physically difficult for the radical to react with another molecule. Examples of these include Gomberg's triphenylmethyl radical, Fremy's salt, such as TEMPO, TEMPOL, nitronyl nitroxides, azephenylenyls and radicals derived from PTM and TTM.
Persistent radicals are generated in great quantity during combustion, "may be responsible for the oxidative stress resulting in cardiopulmonary disease and cancer, attributed to exposure to airborne fine particles". Gomberg's free radical can be generated by following reaction in lab - 3C-Cl + Ag === 3C• + AgCl The reason for persistivity of free radicals is either the delocalisation of unpaired electron or the unavailability of unpaired electron to other species due to the screening of neighbouring atoms/groups. Diradicals are molecules containing two radical centers. Multiple radical centers can exist in a molecule. Atmospheric oxygen exists as a diradical in its ground state as triplet oxygen; the low reactivity of atmospheric oxygen is due to its diradical state. Non-radical states of dioxygen are less stable tha
In organic chemistry, a carbyne is a general term for any compound whose molecular structure includes an electrically neutral carbon atom with three non-bonded electrons, connected to another atom by a single bond. A carbyne thus has the general formula R-C3•, where R is any monovalent group and the superscript 3• indicates the three unbonded electrons. Carbynes are named after the simplest such compound, HC3•, the methylidyne radical or unsubstituted carbyne. Carbyne molecules are found to be in electronic doublet states: the non-bonding electrons on carbon are arranged as one radical and one electron pair, leaving a vacant atomic orbital, rather than being a tri-radical; the simplest case is the CH radical, which has an electron configuration 1σ2 2σ2 3σ2 1π. Here the 1σ molecular orbital is the carbon 1s atomic orbital, the 2σ is the C-H bonding orbital formed by overlap of a carbon s-p hybrid orbital with the hydrogen 1s orbital; the 3σ is a carbon non-bonding orbital pointing along the C-H axis away from the hydrogen, while there are two non-bonding 1π orbitals perpendicular to the C-H axis.
However the 3σ is an s-p hybrid which has lower energy than the 1π orbital, pure p, so the 3σ is filled before the 1π. The CH radical is in fact isoelectronic with the nitrogen atom which does have three unpaired electrons in accordance with Hund's rule of maximum multiplicity; however the nitrogen atom has three degenerate p orbitals, in contrast to the CH radical where hybridization of one orbital leads to an energy difference. A carbyne can occur as a short-lived reactive intermediate. For instance, fluoromethylidyne can be detected in the gas phase by spectroscopy as an intermediate in the flash photolysis of CHFBr2. Carbynes can act as trivalent ligands in complexes with transition metals, in which they are connected to a metal by the three non-bonded electrons in the -C3• group. Examples of such coordination compounds are Cl4W≡C-CH3, WBr2≡C-aryl and WBr22≡C-NR2; such a compound can be obtained by the reaction of tungsten hexacarbonyl W6 with lithium diisopropylamide to form C=W5. This salt is oxidized with either oxalyl bromide or triphenylphosphine dibromide, followed by addition of triphenylphosphine.
Another method is to treat a methoxy metal carbene with a Lewis acid
In quantum mechanics and particle physics, spin is an intrinsic form of angular momentum carried by elementary particles, composite particles, atomic nuclei. Spin is one of two types of angular momentum in quantum mechanics, the other being orbital angular momentum; the orbital angular momentum operator is the quantum-mechanical counterpart to the classical angular momentum of orbital revolution and appears when there is periodic structure to its wavefunction as the angle varies. The existence of spin angular momentum is inferred from experiments, such as the Stern–Gerlach experiment, in which silver atoms were observed to possess two possible discrete angular momenta despite having no orbital angular momentum. In some ways, spin is like a vector quantity. All elementary particles of a given kind have the same magnitude of spin angular momentum, indicated by assigning the particle a spin quantum number; the SI unit of spin is the or, just as with classical angular momentum. In practice, spin is given as a dimensionless spin quantum number by dividing the spin angular momentum by the reduced Planck constant ħ, which has the same units of angular momentum, although this is not the full computation of this value.
The "spin quantum number" is called "spin", leaving its meaning as the unitless "spin quantum number" to be inferred from context. When combined with the spin-statistics theorem, the spin of electrons results in the Pauli exclusion principle, which in turn underlies the periodic table of chemical elements. Wolfgang Pauli in 1924 was the first to propose a doubling of electron states due to a two-valued non-classical "hidden rotation". In 1925, George Uhlenbeck and Samuel Goudsmit at Leiden University suggested the simple physical interpretation of a particle spinning around its own axis, in the spirit of the old quantum theory of Bohr and Sommerfeld. Ralph Kronig anticipated the Uhlenbeck-Goudsmit model in discussion with Hendrik Kramers several months earlier in Copenhagen, but did not publish; the mathematical theory was worked out in depth by Pauli in 1927. When Paul Dirac derived his relativistic quantum mechanics in 1928, electron spin was an essential part of it; as the name suggests, spin was conceived as the rotation of a particle around some axis.
This picture is correct so far as spin obeys the same mathematical laws as quantized angular momenta do. On the other hand, spin has some peculiar properties that distinguish it from orbital angular momenta: Spin quantum numbers may take half-integer values. Although the direction of its spin can be changed, an elementary particle cannot be made to spin faster or slower; the spin of a charged particle is associated with a magnetic dipole moment with a g-factor differing from 1. This could only occur classically if the internal charge of the particle were distributed differently from its mass; the conventional definition of the spin quantum number, s, is s = n/2, where n can be any non-negative integer. Hence the allowed values of s are 1/2, 1, 3/2, 2, etc.. The value of s for an elementary particle depends only on the type of particle, cannot be altered in any known way; the spin angular momentum, S, of any physical system is quantized. The allowed values of S are S = ℏ s = h 4 π n, where h is the Planck constant and ℏ = h/2π is the reduced Planck constant.
In contrast, orbital angular momentum can only take on integer values of s. Those particles with half-integer spins, such as 1/2, 3/2, 5/2, are known as fermions, while those particles with integer spins, such as 0, 1, 2, are known as bosons; the two families of particles obey different rules and broadly have different roles in the world around us. A key distinction between the two families is. In contrast, bosons obey the rules of Bose–Einstein statistics and have no such restriction, so they may "bunch together" if in identical states. Composite particles can have spins different from their component particles. For example, a helium atom in the ground state has spin 0 and behaves like a boson though the quarks and electrons which make it up are all fermions; this has some profound consequences: Quarks and leptons, which make up what is classically known as matter, are all fermions with spin 1/2. The common idea that "matter takes up space" comes from the Pauli exclusion principle acting on these particles to prevent the fermions that make up matter from being in the same quantum state.
Further compaction would require electrons to occupy the same energy states, therefore a kind of pressure acts to resist the fermions being overly close. Elementary fermions with other spins are not known to exist. Elementary particles which are thought of as carrying forces are all bosons with spin 1, they include the photon which carries the electromagnetic force, the gluon, the W and Z bosons. The ability of bosons to occupy the same quantu
Diazomethane is the chemical compound CH2N2, discovered by German chemist Hans von Pechmann in 1894. It is the simplest diazo compound. In the pure form at room temperature, it is an sensitive explosive yellow gas; the compound is a popular methylating agent in the laboratory, but it is too hazardous to be employed on an industrial scale without special precautions. Use of diazomethane has been reduced by the introduction of the safer and equivalent reagent trimethylsilyldiazomethane. For safety and convenience diazomethane is always prepared as needed as a solution in ether and used as such, it converts carboxylic acids into their methyl esters. The reaction is thought to proceed via proton transfer from carboxylic acid to diazomethane to give methyldiazonium cation, which reacts with the carboxylate ion to give the methyl ester and nitrogen gas. Since proton transfer is required and rate limiting, this reaction exhibits high specificity for carboxylic acids over less acidic oxygenated functional groups like alcohols and phenols.
In more specialized applications and homologues are used in Arndt-Eistert synthesis and the Büchner–Curtius–Schlotterbeck reaction for homologation. Diazomethane reacts with phenols in presence of boron trifluoride to give methyl ethers. Diazomethane is frequently used as a carbene source, it takes part in 1,3-dipolar cycloadditions. Diazomethane is prepared by hydrolysis of an ethereal solution of an N-methyl nitrosamide with aqueous base; the traditional precursor is N-nitroso-N-methylurea, but this compound is itself somewhat unstable, nowadays compounds such as N-methyl-N'-nitro-N-nitrosoguanidine and N-methyl-N-nitroso-p-toluenesulfonamide are preferred. CH2N2 reacts with basic solutions of D2O to give the deuterated derivative CD2N2; the concentration of CH2N2 can be determined in either of two convenient ways. It can be treated with an excess of benzoic acid in cold Et2O. Unreacted benzoic acid is back-titrated with standard NaOH. Alternatively, the concentration of CH2N2 in Et2O can be determined spectrophotometrically at 410 nm where its extinction coefficient, ε, is 7.2.
The gas-phase concentration of diazomethane can be determined using photoacoustic spectroscopy. Diazomethane is both isomeric and isoelectronic with the more stable cyanamide, but they cannot interconvert. Many substituted derivatives of diazomethane have been prepared: The stable 2CN2, Ph2CN2. 3SiCHN2, commercially available as a solution and is as effective as CH2N2 for methylation. PhCN2, a red liquid b.p.< 25 °C at 0.1 mm Hg. Diazomethane is toxic by contact with the skin or eyes. Symptoms include chest discomfort, weakness and, in severe cases, collapse. Symptoms may be delayed. Deaths from diazomethane poisoning have been reported. In one instance a laboratory worker consumed a hamburger near a fumehood where he was generating a large quantity of diazomethane, died four days from fulminating pneumonia. Like any other alkylating agent it is expected to be carcinogenic, but such concerns are overshadowed by its serious acute toxicity. CH2N2 may explode in contact with sharp edges, such as ground-glass joints scratches in glassware.
Glassware should be inspected before preparation should take place behind a blast shield. Specialized kits to prepare diazomethane with flame-polished joints are commercially available; the compound explodes when heated beyond 100 °C, exposed to intense light, alkali metals, or calcium sulfate. Use of a blast shield is recommended while using this compound. Proof-of-concept work has been done with microfluidics, in which continuous point-of-use synthesis from N-methyl-N-nitrosourea and 0.93M potassium hydroxide in water was followed by point-of-use conversion with benzoic acid, resulting in a 65% yield of the methyl benzoate ester within seconds at temperatures ranging from 0-50 C. The yield was better than under capillary conditions; the stable compound cyanamide, whose minor tautomer is carbodiimide, is an isomer of diazomethane. Less stable but still isolable isomers of diazomethane include the cyclic 3H-diazirine and isocyanoamine. In addition, the parent nitrilimine has been observed under matrix isolation conditions.
MSDS diazomethane CDC - NIOSH Pocket Guide to Chemical Hazards Sigmaaldrich technical bulletin Sigma-Aldrich diazomethane applications and commercial availability of precursor The Buchner–Curtius–Schlotterbeck reaction @ Institute of Chemistry, Macedonia Identification of Artifacts in Diazomethane and Trimethylsilyldiazomethane Reactions