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
Fullerene chemistry is a field of organic chemistry devoted to the chemical properties of fullerenes. Research in this field tune their properties. For example, fullerene is notoriously insoluble and adding a suitable group can enhance solubility. By adding a polymerizable group, a fullerene polymer can be obtained. Functionalized fullerenes are divided into two classes: exohedral fullerenes with substituents outside the cage and endohedral fullerenes with trapped molecules inside the cage; this article covers the chemistry of these so-called "buckyballs," while the chemistry of carbon nanotubes is covered in carbon nanotube chemistry. Fullerene or C60 is Ih with 12 pentagons and 20 hexagons. According to Euler's theorem these 12 pentagons are required for closure of the carbon network consisting of n hexagons and C60 is the first stable fullerene because it is the smallest possible to obey this rule. In this structure none of the pentagons make contact with each other. Both C60 and its relative C70 obey this so-called isolated pentagon rule.
The next homologue C84 has 24 IPR isomers of which several are isolated and another 51,568 non-IPR isomers. Non-IPR fullerenes have thus far only been isolated as endohedral fullerenes such as Tb3N@C84 with two fused pentagons at the apex of an egg-shaped cage. or as fullerenes with exohedral stabilization such as C50Cl10 and C60H8. Fullerenes with less than 60 carbons do not obey isolated pentagon rule; because of the molecule's spherical shape the carbon atoms are pyramidalized, which has far-reaching consequences for reactivity. It is estimated; the conjugated carbon atoms respond to deviation from planarity by orbital rehybridization of the sp² orbitals and π orbitals to a sp2.27 orbital with a gain in p-character. The p lobes extend further outside the surface than they do into the interior of the sphere and this is one of the reasons a fullerene is electronegative; the other reason is that the empty low-lying π* orbitals have a high s character. The double bonds in fullerene are not all the same.
Two groups can be identified: 30 so-called double bonds connect two hexagons and 60 bonds connect a hexagon and a pentagon. Of the two the bonds are shorter with more double-bond character and therefore a hexagon is represented as a cyclohexatriene and a pentagon as a pentalene or radialene. In other words, although the carbon atoms in fullerene are all conjugated the superstructure is not a super aromatic compound; the X-ray diffraction bond length values are 145.5 pm for the bond. C60 fullerene has 60 π electrons but a closed shell configuration requires 72 electrons; the fullerene is able to acquire the missing electrons by reaction with potassium to form first the K6C6−60 salt and the K12C12−60 In this compound the bond length alternation observed in the parent molecule has vanished. Fullerenes tend to react as electrophiles. An additional driving force is relief of strain. Key in this type of reaction is the level of functionalization i.e. monoaddition or multiple additions and in case of multiple additions their topological relationships.
In conformity with IUPAC rules, the terms methanofullerene are used to indicate the ring-closed fullerene derivatives, fulleroid to ring-open structures. Fullerenes react as electrophiles with a host of nucleophiles in nucleophilic additions; the intermediary formed carbanion is captured by another electrophile. Examples of nucleophiles are Grignard reagents and organolithium reagents. For example, the reaction of C60 with methylmagnesium chloride stops quantitatively at the penta-adduct with the methyl groups centered around a cyclopentadienyl anion, subsequently protonated. Another nucleophilic reaction is the Bingel reaction. Fullerene reacts with chlorobenzene and aluminium chloride in a Friedel-Crafts alkylation type reaction. In this hydroarylation the reaction product is the 1,2-addition adduct; the bonds of fullerenes react as dienes or dienophiles in cycloadditions for instance Diels-Alder reactions. 4-membered rings can be obtained by cycloadditions for instance with benzyne. An example of a 1,3-dipolar cycloaddition to a 5-membered ring is the Prato reaction.
Fullerenes are hydrogenated by several methods. Examples of hydrofullerenes are C60H18 and C60H36; however hydrogenated C60H60 is only hypothetical because of large strain. Hydrogenated fullerenes are not stable, as prolonged hydrogenation of fullerenes by direct reaction with hydrogen gas at high temperature conditions results in cage fragmentation. At the final reaction stage this causes collapse of cage structure with formation of polycyclic aromatic hydrocarbons. Although more difficult than reduction, oxidation of fullerene is possible for instance with oxygen and osmium tetraoxide. Fullerenes can be hydroxylated to fullerols. Water solubility depends on the total number of hydroxyl groups. One method is fullerene reaction in diluted sulfuric acid and potassium nitrate to C6015. Another method is reaction in diluted sodium hydroxide catalysed by TBAH adding 24 to 26 hydroxyl groups. Hydroxylation has been reported using solvent-free NaOH / hydrogen peroxide. C608 was prepared using a multistep procedure starting from a mixed peroxide fullerene.
The maximum number of hydroxyl groups that can be attached stands at 36–40. Fullerenes react in electrophilic additions as well; the reaction with bromine can add up to 24 bromine atoms to the sphere. The
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
Sodium hydride is the chemical compound with the empirical formula NaH. This alkali metal hydride is used as a strong, yet combustible base in organic synthesis. NaH is representative of the saline hydrides, meaning it is a salt-like hydride, composed of Na+ and H− ions, in contrast to the more molecular hydrides such as borane, methane and water, it is an ionic material, insoluble in organic solvents, consistent with the fact that H− remains an unknown anion in solution. Because of the insolubility of NaH, all reactions involving NaH occur at the surface of the solid. NaH is produced by the direct reaction of liquid sodium. Pure NaH is colorless, although samples appear grey. NaH is ca. 40% denser than Na. NaH, like LiH, KH, RbH, CsH, adopts the NaCl crystal structure. In this motif, each Na+ ion is surrounded by six H− centers in an octahedral geometry; the ionic radii of H − and F − are comparable, as judged by the Na − Na − F distances. A unusual situation occurs in a compound dubbed "inverse sodium hydride", which contains Na− and H+ ions.
Na− is an alkalide, this compound differs from ordinary sodium hydride in having a much higher energy content due to the net displacement of two electrons from hydrogen to sodium. A derivative of this "inverse sodium hydride" arises in the presence of the base adamanzane; this molecule irreversibly encapsulates the H+ and shields it from interaction with the alkalide Na−. Theoretical work has suggested that an unprotected protonated tertiary amine complexed with the sodium alkalide might be metastable under certain solvent conditions, though the barrier to reaction would be small and finding a suitable solvent might be difficult. NaH is a base of wide utility in organic chemistry; as a superbase, it is capable of deprotonating a range of weak Brønsted acids to give the corresponding sodium derivatives. Typical "easy" substrates contain O-H, N-H, S-H bonds, including alcohols, phenols and thiols. NaH most notably is employed to deprotonate carbon acids such as 1,3-dicarbonyls and analogues such as malonic esters.
The resulting sodium derivatives can be alkylated. NaH is used to promote condensation reactions of carbonyl compounds via the Dieckmann condensation, Stobbe condensation, Darzens condensation, Claisen condensation. Other carbon acids susceptible to deprotonation by NaH include sulfonium salts and DMSO. NaH is used to make sulfur ylides, which in turn are used to convert ketones into epoxides, as in the Johnson–Corey–Chaykovsky reaction. NaH reduces certain main group compounds, but analogous reactivity is rare in organic chemistry. Notably boron trifluoride reacts to give diborane and sodium fluoride: 6 NaH + 2 BF3 → B2H6 + 6 NaFSi-Si and S-S bonds in disilanes and disulfides are reduced. A series of reduction reactions, including the hydrodecyanation of tertiary nitriles, reduction of imines to amines, amides to aldehydes, can be effected by a composite reagent composed of sodium hydride and an alkali metal iodide; the use of sodium hydride has been proposed for hydrogen storage for use in fuel cell vehicles, the hydride being encased in plastic pellets which are crushed in the presence of water to release the hydrogen.
Sodium hydride is sold by many chemical suppliers as a mixture of 60% sodium hydride in mineral oil. Such a dispersion is safer to handle and weigh than pure NaH; the compound is used in this form but the pure grey solid can be prepared by rinsing the oil with pentane or THF, with care being taken because the washings will contain traces of NaH that can ignite in air. Reactions involving NaH require an inert atmosphere, such as argon gas. NaH is used as a suspension in THF, a solvent that resists deprotonation but solvates many organosodium compounds. NaH can ignite in air upon contact with water to release hydrogen, flammable. Hydrolysis converts NaH into a caustic base. In practice, most sodium hydride is dispensed as a dispersion in oil, which can be safely handled in air
Electrochemistry is the branch of physical chemistry that studies the relationship between electricity, as a measurable and quantitative phenomenon, identifiable chemical change, with either electricity considered an outcome of a particular chemical change or vice versa. These reactions involve electric charges moving between an electrolyte, thus electrochemistry deals with the interaction between electrical energy and chemical change. When a chemical reaction is caused by an externally supplied current, as in electrolysis, or if an electric current is produced by a spontaneous chemical reaction as in a battery, it is called an electrochemical reaction. Chemical reactions where electrons are transferred directly between molecules and/or atoms are called oxidation-reduction or reactions. In general, electrochemistry describes the overall reactions when individual redox reactions are separate but connected by an external electric circuit and an intervening electrolyte. Understanding of electrical matters began in the sixteenth century.
During this century, the English scientist William Gilbert spent 17 years experimenting with magnetism and, to a lesser extent, electricity. For his work on magnets, Gilbert became known as the "Father of Magnetism." He discovered various methods for strengthening magnets. In 1663, the German physicist Otto von Guericke created the first electric generator, which produced static electricity by applying friction in the machine; the generator was made of a large sulfur ball cast inside a glass globe, mounted on a shaft. The ball was rotated by means of a crank and an electric spark was produced when a pad was rubbed against the ball as it rotated; the globe could be used as source for experiments with electricity. By the mid—18th century the French chemist Charles François de Cisternay du Fay had discovered two types of static electricity, that like charges repel each other whilst unlike charges attract. Du Fay announced that electricity consisted of two fluids: positive, electricity; this was the two-fluid theory of electricity, to be opposed by Benjamin Franklin's one-fluid theory in the century.
In 1785, Charles-Augustin de Coulomb developed the law of electrostatic attraction as an outgrowth of his attempt to investigate the law of electrical repulsions as stated by Joseph Priestley in England. In the late 18th century the Italian physician and anatomist Luigi Galvani marked the birth of electrochemistry by establishing a bridge between chemical reactions and electricity on his essay "De Viribus Electricitatis in Motu Musculari Commentarius" in 1791 where he proposed a "nerveo-electrical substance" on biological life forms. In his essay Galvani concluded that animal tissue contained a here-to-fore neglected innate, vital force, which he termed "animal electricity," which activated nerves and muscles spanned by metal probes, he believed that this new force was a form of electricity in addition to the "natural" form produced by lightning or by the electric eel and torpedo ray as well as the "artificial" form produced by friction. Galvani's scientific colleagues accepted his views, but Alessandro Volta rejected the idea of an "animal electric fluid," replying that the frog's legs responded to differences in metal temper and bulk.
Galvani refuted this by obtaining muscular action with two pieces of the same material. In 1800, William Nicholson and Johann Wilhelm Ritter succeeded in decomposing water into hydrogen and oxygen by electrolysis. Soon thereafter Ritter discovered the process of electroplating, he observed that the amount of metal deposited and the amount of oxygen produced during an electrolytic process depended on the distance between the electrodes. By 1801, Ritter observed thermoelectric currents and anticipated the discovery of thermoelectricity by Thomas Johann Seebeck. By the 1810s, William Hyde Wollaston made improvements to the galvanic cell. Sir Humphry Davy's work with electrolysis led to the conclusion that the production of electricity in simple electrolytic cells resulted from chemical action and that chemical combination occurred between substances of opposite charge; this work led directly to the isolation of sodium and potassium from their compounds and of the alkaline earth metals from theirs in 1808.
Hans Christian Ørsted's discovery of the magnetic effect of electric currents in 1820 was recognized as an epoch-making advance, although he left further work on electromagnetism to others. André-Marie Ampère repeated Ørsted's experiment, formulated them mathematically. In 1821, Estonian-German physicist Thomas Johann Seebeck demonstrated the electrical potential in the juncture points of two dissimilar metals when there is a heat difference between the joints. In 1827, the German scientist Georg Ohm expressed his law in this famous book "Die galvanische Kette, mathematisch bearbeitet" in which he gave his complete theory of electricity. In 1832, Michael Faraday's experiments led him to state his two laws of electrochemistry. In 1836, John Daniell invented a primary cell which solved the problem of polarization by eliminating hydrogen gas generation at the positive electrode. Results revealed that alloying the amalgamated zinc with mercury would produce a higher voltage. William Grove produced the first fuel cell in 1839.
In 1846, Wilhelm Weber developed the electrodynamometer. In 1868, Georges Leclanché patented a new cell which became the forerunner to the world's first used battery, the zinc carbon cell. Svante Arrhenius published
In chemistry, a leaving group is a molecular fragment that departs with a pair of electrons in heterolytic bond cleavage. Leaving groups can be anions or neutral molecules, but in either case it is crucial that the leaving group be able to stabilize the additional electron density that results from bond heterolysis. Common anionic leaving groups are halides such as Cl−, Br−, I−, sulfonate esters such as tosylate. Fluoride functions as a leaving group in the nerve-agent sarin gas. Common neutral molecule leaving groups are ammonia. Leaving groups may be positively charged cations; the physical manifestation of leaving group ability is the rate. Good leaving groups give fast reactions. By transition state theory, this implies that reactions involving good leaving groups have low activation barriers leading to stable transition states, it is helpful to consider the concept of leaving group ability in the case of the first step of an SN1/E1 reaction with an anionic leaving group, while keeping in mind that this concept can be generalized to all reactions that involve leaving groups.
Because the leaving group bears a larger negative charge in the transition state than in the starting material, a good leaving group must be able to stabilize this negative charge, i.e. form stable anions. A good measure of anion stability is the pKa of an anion's conjugate acid, leaving group ability indeed follows this trend, with a lower pKaH being associated with better leaving group ability; the correlation between pKaH and leaving group ability, however, is not perfect. Leaving group ability represents the difference in energy between starting materials and a transition state, differences in leaving group ability are reflected in changes in this quantity; the quantity pKaH, represents the difference in energy between starting materials and products with differences in acidity reflected in changes in this quantity. The starting materials in these cases are different. In the case of pKa, the "leaving group" is bound to a proton in the starting material, while in the case of leaving group ability, the leaving group is bound to carbon.
It is with these important caveats in mind that one must consider pKaH to be reflective of leaving group ability, but the trends in each tend to correlate well with each other. Consistent with this picture, strong bases such as OH−, OR− and NR2− tend to make poor leaving groups, due their inability to stabilize a negative charge, it is exceedingly rare for groups such as H− and R3C− to depart with a pair of electrons because of the instability of these bases. It is important to note that the list given above describes trends; the ability of a group to leave is contextual. For example, in SNAr reactions, the rate is increased when the leaving group is fluoride relative to the other halogens; this effect is due to the fact that the highest energy transition state for this two step addition-elimination process occurs in the first step, where fluoride's greater electron withdrawing capability relative to the other halides stabilizes the developing negative charge on the aromatic ring. The departure of the leaving group takes place from this high energy Meisenheimer complex, since the departure is not involved in the rate limiting step, it does not affect the overall rate of the reaction.
This effect is general to conjugate base eliminations. When the departure of the leaving group is involved in the rate limiting step of a reaction there can still exist contextual differences that can change the order of leaving group ability. In Friedel-Crafts alkylations, the normal halogen leaving group order is reversed so that the rate of the reaction follows RF > RCl > RBr > RI. This effect is due to their greater ability to complex the Lewis acid catalyst, the actual group that leaves is an "ate" complex between the Lewis acid and the departing leaving group; this situation is broadly defined as leaving group activation There can still exist contextual differences in leaving group ability in the purest form, when the actual group that leaves is not affected by the reaction conditions and the departure of the leaving group occurs in the rate determining step. In the situation where other variables are held constant, a change in nucleophile can lead to a change in the order of reactivity for leaving groups.
In the case below, tosylate is the best leaving group when ethoxide is the nucleophile, but iodide and bromide become better leaving groups in the case of the thiolate nucleophile. It is common in E1 and SN1 reactions for a poor leaving group to be transformed into a good one by protonation or complexation with a Lewis acid. Thus, it is by protonation prior to departure that a molecule can formally lose such poor leaving groups as hydroxide; the same principle is at work in the Friedel-Crafts reaction. Here, a strong Lewis acid is required to generate either a carbocation from an alkyl halide in the Friedel-Crafts alkylation reaction or an acylium ion from an acyl halide. In the vast majority of cases, reactions that involve leaving group activation generate a cation in a separate step, prior to either nucleophilic attack or elimination. For example, SN1 and E1 reactions may involve an activation step, whereas SN2 and E2 reactions do not; the requirement for a good leaving group is relaxed in conjugate base elimination reactions.
These reactions include loss of a leaving group in the β position of an eno