Lewis acids and bases
A Lewis acid is a chemical species that contains an empty orbital, capable of accepting an electron pair from a Lewis base to form a Lewis adduct. A Lewis base is any species that has a filled orbital containing an electron pair, not involved in bonding but may form a dative bond with a Lewis acid to form a Lewis adduct. For example, NH3 is a Lewis base. Trimethylborane is a Lewis acid. In a Lewis adduct, the Lewis acid and base share an electron pair furnished by the Lewis base, forming a dative bond. In the context of a specific chemical reaction between NH3 and Me3B, the lone pair from NH3 will form a dative bond with the empty orbital of Me3B to form an adduct NH3•BMe3; the terminology refers to the contributions of Gilbert N. Lewis; the terms nucleophile and electrophile are more or less interchangeable with Lewis base and Lewis acid, respectively. However, these terms their abstract noun forms nucleophilicity and electrophilicity, emphasize the kinetic aspect of reactivity, while the Lewis basicity and Lewis acidity emphasize the thermodynamic aspect of Lewis adduct formation.
In many cases, the interaction between the Lewis base and Lewis acid in a complex is indicated by an arrow indicating the Lewis base donating electrons toward the Lewis acid using the notation of a dative bond—for example, Me3B←NH3. Some sources indicate the Lewis base with a pair of dots, which allows consistent representation of the transition from the base itself to the complex with the acid: Me3B +:NH3 → Me3B:NH3A center dot may be used to represent a Lewis adduct, such as Me3B•NH3. Another example is boron trifluoride diethyl etherate, BF3•Et2O. Although there have been attempts to use computational and experimental energetic criteria to distinguish dative bonding from non-dative covalent bonds, for the most part, the distinction makes note of the source of the electron pair, dative bonds, once formed, behave as other covalent bonds do, though they have considerable polar character. Moreover, in some cases, the use of the dative bond arrow is just a notational convenience for avoiding the drawing of formal charges.
In general, the donor–acceptor bond is viewed as somewhere along a continuum between idealized covalent bonding and ionic bonding. Classically, the term "Lewis acid" is restricted to trigonal planar species with an empty p orbital, such as BR3 where R can be an organic substituent or a halide. For the purposes of discussion complex compounds such as Et3Al2Cl3 and AlCl3 are treated as trigonal planar Lewis acids. Metal ions such as Na+, Mg2+, Ce3+, which are invariably complexed with additional ligands, are sources of coordinatively unsaturated derivatives that form Lewis adducts upon reaction with a Lewis base. Other reactions might be referred to as "acid-catalyzed" reactions; some compounds, such as H2O, are both Lewis acids and Lewis bases, because they can either accept a pair of electrons or donate a pair of electrons, depending upon the reaction. Lewis acids are diverse. Simplest are those, but more common are those. Examples of Lewis acids based on the general definition of electron pair acceptor include: the proton and acidic compounds onium ions, such as NH4+ and H3O+ high oxidation state transition metal cations, e.g. Fe3+.
Again, the description of a Lewis acid is used loosely. For example, in solution, bare protons do not exist; some of the most studied examples of such Lewis acids are the boron trihalides and organoboranes, but other compounds exhibit this behavior: BF3 + F− → BF4−In this adduct, all four fluoride centres are equivalent. BF3 + OMe2 → BF3OMe2Both BF4− and BF3OMe2 are Lewis base adducts of boron trifluoride. In many cases, the adducts violate the octet rule, such as the triiodide anion: I2 + I− → I3−The variability of the colors of iodine solutions reflects the variable abilities of the solvent to form adducts with the Lewis acid I2. In some cases, the Lewis acid is capable of binding two Lewis base, a famous example being the formation of hexafluorosilicate: SiF4 + 2 F− → SiF62− Most compounds considered to be Lewis acids require an activation step prior to formation of the adduct with the Lewis base. Well known cases are the aluminium trihalides, which are viewed as Lewis acids. Aluminium trihalides, unlike the boron trihalides, do not exist in the form AlX3, but as aggregates and polymers that must be degraded by the Lewis base.
A simpler case is the formation of adducts of borane. Monomeric BH3 does not exist appreciably, so the adducts of borane are generated by degradation of diborane: B2H6 + 2 H− → 2 BH4−In this case, an intermediate B2H7− can be isolated. Many metal complexes serve as Lewis acids, but only after dissociating a more weakly bound Lewis base water. 2+ + 6 NH3 → 2+ + 6 H2O The proton is one of the strongest but is one of the most complicated Lewis acids. It is convention to ignore the fact that a proton is solvated (bound to solvent
In organic chemistry, the term aromaticity is used to describe a cyclic, planar molecule with a ring of resonance bonds that exhibits more stability than other geometric or connective arrangements with the same set of atoms. Aromatic molecules are stable, do not break apart to react with other substances. Organic compounds that are not aromatic are classified as aliphatic compounds—they might be cyclic, but only aromatic rings have special stability. Since the most common aromatic compounds are derivatives of benzene, the word aromatic refers informally to benzene derivatives, so it was first defined. Many non-benzene aromatic compounds exist. In living organisms, for example, the most common aromatic rings are the double-ringed bases in RNA and DNA. An aromatic functional group or other substituent is called an aryl group; the earliest use of the term aromatic was in an article by August Wilhelm Hofmann in 1855. Hofmann used the term for a class of benzene compounds, many of which have odors, unlike pure saturated hydrocarbons.
Aromaticity as a chemical property bears no general relationship with the olfactory properties of such compounds, although in 1855, before the structure of benzene or organic compounds was understood, chemists like Hofmann were beginning to understand that odiferous molecules from plants, such as terpenes, had chemical properties that we recognize today are similar to unsaturated petroleum hydrocarbons like benzene. In terms of the electronic nature of the molecule, aromaticity describes a conjugated system made of alternating single and double bonds in a ring; this configuration allows for the electrons in the molecule's pi system to be delocalized around the ring, increasing the molecule's stability. The molecule cannot be represented by one structure, but rather a resonance hybrid of different structures, such as with the two resonance structures of benzene; these molecules cannot be found in either one of these representations, with the longer single bonds in one location and the shorter double bond in another.
Rather, the molecule exhibits bond lengths in between those of double bonds. This seen model of aromatic rings, namely the idea that benzene was formed from a six-membered carbon ring with alternating single and double bonds, was developed by August Kekulé; the model for benzene consists of two resonance forms, which corresponds to the double and single bonds superimposing to produce six one-and-a-half bonds. Benzene is a more stable molecule than would be expected without accounting for charge delocalization; as it is a standard for resonance diagrams, the use of a double-headed arrow indicates that two structures are not distinct entities but hypothetical possibilities. Neither is an accurate representation of the actual compound, best represented by a hybrid of these structures. A C=C bond is shorter than a C−C bond. Benzene is a regular hexagon—it is planar and all six carbon–carbon bonds have the same length, intermediate between that of a single and that of a double bond. In a cyclic molecule with three alternating double bonds, the bond length of the single bond would be 1.54 Å and that of the double bond would be 1.34 Å.
However, in a molecule of benzene, the length of each of the bonds is 1.40 Å, indicating it to be the average of single and double bond. A better representation is that of the circular π-bond, in which the electron density is evenly distributed through a π-bond above and below the ring; this model more represents the location of electron density within the aromatic ring. The single bonds are formed from overlap of hybridized atomic sp2-orbitals in line between the carbon nuclei—these are called σ-bonds. Double bonds consist of a π-bond; the π-bonds are formed from overlap of atomic p-orbitals below the plane of the ring. The following diagram shows the positions of these p-orbitals: Since they are out of the plane of the atoms, these orbitals can interact with each other and become delocalized; this means that, instead of being tied to one atom of carbon, each electron is shared by all six in the ring. Thus, there are not enough electrons to form double bonds on all the carbon atoms, but the "extra" electrons strengthen all of the bonds on the ring equally.
The resulting molecular orbital is considered to have π symmetry. The first known use of the word "aromatic" as a chemical term—namely, to apply to compounds that contain the phenyl group—occurs in an article by August Wilhelm Hofmann in 1855. If this is indeed the earliest introduction of the term, it is curious that Hofmann says nothing about why he introduced an adjective indicating olfactory character to apply to a group of chemical substances only some of which have notable aromas. Many of the most odoriferous organic substances known are terpenes, which are not aromatic in the chemical sense, but terpenes and benzenoid substances do have a chemical characteristic in common, namely higher unsaturation than many aliphatic compounds, Hofmann may not have been making a distinction between the two categories. Many of the earliest-known examples of aromatic compounds, such as benzene and toluene, have distinctive pleasant smells; this property led to the term "aromatic" for this class of compounds, hence the term "aromaticity" for the discovered electronic property.
In the 19th century chemists found it puzzling that benzene could be so unreactive toward addition reactions, given its presumed high degree of unsaturation. The cyclohexatriene structure for benzene was first pr
Sulfuric acid known as vitriol, is a mineral acid composed of the elements sulfur and hydrogen, with molecular formula H2SO4. It is a colorless and syrupy liquid, soluble in water, in a reaction, exothermic, its corrosiveness can be ascribed to its strong acidic nature, and, if at a high concentration, its dehydrating and oxidizing properties. It is hygroscopic absorbing water vapor from the air. Upon contact, sulfuric acid can cause severe chemical burns and secondary thermal burns. Sulfuric acid is a important commodity chemical, a nation's sulfuric acid production is a good indicator of its industrial strength, it is produced with different methods, such as contact process, wet sulfuric acid process, lead chamber process and some other methods. Sulfuric acid is a key substance in the chemical industry, it is most used in fertilizer manufacture, but is important in mineral processing, oil refining, wastewater processing, chemical synthesis. It has a wide range of end applications including in domestic acidic drain cleaners, as an electrolyte in lead-acid batteries, in various cleaning agents.
Although nearly 100% sulfuric acid can be made, the subsequent loss of SO3 at the boiling point brings the concentration to 98.3% acid. The 98.3% grade is more stable in storage, is the usual form of what is described as "concentrated sulfuric acid". Other concentrations are used for different purposes; some common concentrations are: "Chamber acid" and "tower acid" were the two concentrations of sulfuric acid produced by the lead chamber process, chamber acid being the acid produced in the lead chamber itself and tower acid being the acid recovered from the bottom of the Glover tower. They are now obsolete as commercial concentrations of sulfuric acid, although they may be prepared in the laboratory from concentrated sulfuric acid if needed. In particular, "10M" sulfuric acid is prepared by adding 98% sulfuric acid to an equal volume of water, with good stirring: the temperature of the mixture can rise to 80 °C or higher. Sulfuric acid reacts with its anhydride, SO3, to form H2S2O7, called pyrosulfuric acid, fuming sulfuric acid, Disulfuric acid or oleum or, less Nordhausen acid.
Concentrations of oleum are either expressed in terms of % SO3 or as % H2SO4. Pure H2S2O7 is a solid with melting point of 36 °C. Pure sulfuric acid has a vapor pressure of <0.001 mmHg at 25 °C and 1 mmHg at 145.8 °C, 98% sulfuric acid has a <1 mmHg vapor pressure at 40 °C. Pure sulfuric acid is a viscous clear liquid, like oil, this explains the old name of the acid. Commercial sulfuric acid is sold in several different purity grades. Technical grade H2SO4 is impure and colored, but is suitable for making fertilizer. Pure grades, such as United States Pharmacopeia grade, are used for making pharmaceuticals and dyestuffs. Analytical grades are available. Nine hydrates are known, but four of them were confirmed to be tetrahydrate and octahydrate. Anhydrous H2SO4 is a polar liquid, having a dielectric constant of around 100, it has a high electrical conductivity, caused by dissociation through protonating itself, a process known as autoprotolysis. 2 H2SO4 ⇌ H3SO+4 + HSO−4The equilibrium constant for the autoprotolysis is Kap = = 2.7×10−4The comparable equilibrium constant for water, Kw is 10−14, a factor of 1010 smaller.
In spite of the viscosity of the acid, the effective conductivities of the H3SO+4 and HSO−4 ions are high due to an intramolecular proton-switch mechanism, making sulfuric acid a good conductor of electricity. It is an excellent solvent for many reactions; because the hydration reaction of sulfuric acid is exothermic, dilution should always be performed by adding the acid to the water rather than the water to the acid. Because the reaction is in an equilibrium that favors the rapid protonation of water, addition of acid to the water ensures that the acid is the limiting reagent; this reaction is best thought of as the formation of hydronium ions: H2SO4 + H2O → H3O+ + HSO−4 Ka1 = 2.4×106 HSO−4 + H2O → H3O+ + SO2−4 Ka2 = 1.0×10−2 HSO−4 is the bisulfate anion and SO2−4 is the sulfate anion. Ka1 and Ka2 are the acid dissociation constants; because the hydration of sulfuric acid is thermodynamically favorable and the affinity of it for water is sufficiently strong, sulfuric acid is an excellent dehydrating agent.
Concentrated sulfuric acid has a powerful dehydrating property, removing water from other chemical compounds including sugar and other carbohydrates and producing carbon and steam. In the laboratory, this is demonstrated by mixing table sugar into sulfuric acid; the sugar changes from white to dark brown and to black as carbon is formed. A rigid column of black, porous carbon will emerge as well; the carbon will smell of caramel due to the heat generated. C 12 H 22 O 11 ⏞ sucrose → H 2 SO 4 12 C + 11 H 2
A cyclic compound is a term for a compound in the field of chemistry in which one or more series of atoms in the compound is connected to form a ring. Rings may vary in size from three to many atoms, include examples where all the atoms are carbon, none of the atoms are carbon, or where both carbon and non-carbon atoms are present. Depending on the ring size, the bond order of the individual links between ring atoms, their arrangements within the rings and heterocyclic compounds may be aromatic or non-aromatic, in the latter case, they may vary from being saturated to having varying numbers of multiple bonds between the ring atoms; because of the tremendous diversity allowed, in combination, by the valences of common atoms and their ability to form rings, the number of possible cyclic structures of small size numbers in the many billions. Cyclic compound examples: All-carbon and more complex natural cyclic compounds. Adding to their complexity and number, closing of atoms into rings may lock particular atoms with distinct substitution such that stereochemistry and chirality of the compound results, including some manifestations that are unique to rings.
As well, depending on ring size, the three-dimensional shapes of particular cyclic structures—typically rings of 5-atoms and larger—can vary and interconvert such that conformational isomerism is displayed. Indeed, the development of this important chemical concept arose in reference to cyclic compounds. Cyclic compounds, because of the unique shapes, reactivities and bioactivities that they engender, are the largest majority of all molecules involved in the biochemistry and function of living organisms, in the man-made molecules. A cyclic compound or ring compound is a compound at least some of whose atoms are connected to form a ring. Rings vary in size from 3 to many tens or hundreds of atoms. Examples of ring compounds include cases where: all the atoms are carbon, none of the atoms are carbon, or where both carbon and non-carbon atoms are present. Common atoms can form varying numbers of bonds, many common atoms form rings. In addition, depending on the ring size, the bond order of the individual links between ring atoms, their arrangements within the rings, cyclic compounds may be aromatic or non-aromatic.
As a consequence of the constitutional variability, thermodynamically possible in cyclic structures, the number of possible cyclic structures of small size numbers in the many billions. Moreover, the closing of atoms into rings may lock particular functional group–substituted atoms into place, resulting in stereochemistry and chirality being associated with the compound, including some manifestations that are unique to rings. IUPAC nomenclature has extensive rules to cover the naming of cyclic structures, both as core structures, as substituents appended to alicyclic structures; the term macrocycle is used. The term polycyclic is used. Naphthalene is formally a polycyclic compound, but is more named as a bicyclic compound. Several examples of macrocyclic and polycyclic structures are given in the final gallery below; the atoms that are part of the ring structure are called annular atoms. The vast majority of cyclic compounds are organic, of these, a significant and conceptually important portion are composed of rings made only of carbon atoms.
Inorganic atoms form cyclic compounds as well. Examples include sulfur, silicon and boron; when carbon in benzene is "replaced" by other elements, e.g. as in borabenzene, germanabenzene and phosphorine, aromaticity is retained, so aromatic inorganic cyclic compounds are known and well-characterized. Cyclic compounds that have both carbon and non-carbon atoms present are termed. Hantzsch–Widman nomenclature is recommended by the IUPAC for naming heterocycles, but many common names remain in regular use. Cyclic compounds may not exhibit aromaticity. In organic chemistry, the term aromaticity is used to describe a cyclic, planar molecule that exhibits unusual stability as compared to other geometric or connective arrangements of the same set of atoms; as a result of their stability, it is difficult to cause aromatic molecules to break apart and to react with other substances. Organic compounds that are not aromatic are classified as aliphatic compounds—they might be cyclic, but only aromatic rings have especial stability.
Since one of the most encountered aromatic systems of compounds in organic chemistry is based on derivatives of the prototypic
Maya blue is a unique bright azure blue pigment manufactured by cultures of pre-Columbian Mesoamerica, such as the Maya and Aztec. The Maya blue pigment is a composite of organic and inorganic constituents indigo dyes derived from the leaves of añil plants combined with palygorskite, a natural clay which, mysteriously, is not known to exist in abundant deposits in Mesoamerica. Smaller trace amounts of other mineral additives have been identified. Maya blue first appeared around 800, it was still used in the 16th century in several Convents of Colonial Mexico, notably in the paintings of the Indian Juan Gerson in Tecamachalco; these paintings are a clear example of the combination of Indian and European techniques sometimes known as Arte Indocristiano. After that, the techniques for its production were lost in Mexico, but in Cuba there are examples from as late as 1830. Despite time and the harsh weathering conditions, paintings coloured by Maya blue have not faded over time. More remarkably, the color has resisted chemical acids such as nitric acid.
Its resistance against chemical aggression and biodegradation was tested, it was shown that Maya blue is an resistant pigment, but it can be destroyed using intense acid treatment under reflux. The chemical composition of the compound was determined by powder diffraction in the 1950s and was found to be a composite of palygorskite and indigo, most derived from the leaves of the añil. An actual recipe to reproduce Maya blue pigment was published in 1993 by a Mexican historian and chemist, Constantino Reyes-Valerio; the combination of different clays, together with the use of the leaves of the añil and the actual process is described in his paper. Reyes-Valerio's contributions were due to his combined background of history and chemistry, through a thorough revision of primary texts, microscopic analysis of the mural paintings and fourier transform infrared spectroscopy. After the formula for the production was published in the book De Bonampak al Templo Mayor: Historia del Azul Maya en Mesoamerica, many developments in the chemical analysis of the pigment occurred in collaborations between Reyes-Valerio and European scientists.
A comprehensive study on the pigment which describes history, the experimental study techniques, the syntheses and nature of Maya blue and the research in relation with the archaeological and historical contexts has been published in the journal Developments in Clay Science. Pre-Columbian American culture In the Americas, Maya blue was used as a colorant in pre-Columbian artworks, sculptures and textiles, to illuminate Mesoamerican codices. For example, many illustrations in the Florentine Codex written by Bernardino de Sahagún contain the maya blue color. Maya blue may have been used in the Grolier Codex, yet there's a dispute about this codex, whether or not the use of Maya blue has been confirmed. Recent research suggests Maya blue may have played an important role in human sacrifices to Chaac at Chichén Itzá, both produced at the sacrificial site and used to paint the bodies of the victims. Maya blue is associated with the center of a flame. Holding the most heat and therefore the most tonalli, the blue color is considered precious.
Other classic artificial blue pigments: Egyptian blue, Chinese blue List of colors Azul Maya, descriptive site by Reyes-Valerio
The melting point of a substance is the temperature at which it changes state from solid to liquid. At the melting point the solid and liquid phase exist in equilibrium; the melting point of a substance depends on pressure and is specified at a standard pressure such as 1 atmosphere or 100 kPa. When considered as the temperature of the reverse change from liquid to solid, it is referred to as the freezing point or crystallization point; because of the ability of some substances to supercool, the freezing point is not considered as a characteristic property of a substance. When the "characteristic freezing point" of a substance is determined, in fact the actual methodology is always "the principle of observing the disappearance rather than the formation of ice", that is, the melting point. For most substances and freezing points are equal. For example, the melting point and freezing point of mercury is 234.32 kelvins. However, certain substances possess differing solid-liquid transition temperatures.
For example, agar melts at 85 °C and solidifies from 31 °C. The melting point of ice at 1 atmosphere of pressure is close to 0 °C. In the presence of nucleating substances, the freezing point of water is not always the same as the melting point. In the absence of nucleators water can exist as a supercooled liquid down to −48.3 °C before freezing. The chemical element with the highest melting point is tungsten, at 3,414 °C; the often-cited carbon does not melt at ambient pressure but sublimes at about 3,726.85 °C. Tantalum hafnium carbide is a refractory compound with a high melting point of 4215 K. At the other end of the scale, helium does not freeze at all at normal pressure at temperatures arbitrarily close to absolute zero. Many laboratory techniques exist for the determination of melting points. A Kofler bench is a metal strip with a temperature gradient. Any substance can be placed on a section of the strip, revealing its thermal behaviour at the temperature at that point. Differential scanning calorimetry gives information on melting point together with its enthalpy of fusion.
A basic melting point apparatus for the analysis of crystalline solids consists of an oil bath with a transparent window and a simple magnifier. The several grains of a solid are placed in a thin glass tube and immersed in the oil bath; the oil bath is heated and with the aid of the magnifier melting of the individual crystals at a certain temperature can be observed. In large/small devices, the sample is placed in a heating block, optical detection is automated; the measurement can be made continuously with an operating process. For instance, oil refineries measure the freeze point of diesel fuel online, meaning that the sample is taken from the process and measured automatically; this allows for more frequent measurements as the sample does not have to be manually collected and taken to a remote laboratory. For refractory materials the high melting point may be determined by heating the material in a black body furnace and measuring the black-body temperature with an optical pyrometer. For the highest melting materials, this may require extrapolation by several hundred degrees.
The spectral radiance from an incandescent body is known to be a function of its temperature. An optical pyrometer matches the radiance of a body under study to the radiance of a source, calibrated as a function of temperature. In this way, the measurement of the absolute magnitude of the intensity of radiation is unnecessary. However, known temperatures must be used to determine the calibration of the pyrometer. For temperatures above the calibration range of the source, an extrapolation technique must be employed; this extrapolation is accomplished by using Planck's law of radiation. The constants in this equation are not known with sufficient accuracy, causing errors in the extrapolation to become larger at higher temperatures. However, standard techniques have been developed to perform this extrapolation. Consider the case of using gold as the source. In this technique, the current through the filament of the pyrometer is adjusted until the light intensity of the filament matches that of a black-body at the melting point of gold.
This establishes the primary calibration temperature and can be expressed in terms of current through the pyrometer lamp. With the same current setting, the pyrometer is sighted on another black-body at a higher temperature. An absorbing medium of known transmission is inserted between this black-body; the temperature of the black-body is adjusted until a match exists between its intensity and that of the pyrometer filament. The true higher temperature of the black-body is determined from Planck's Law; the absorbing medium is removed and the current through the filament is adjusted to match the filament intensity to that of the black-body. This establishes a second calibration point for the pyrometer; this step is repeated to carry the calibration to hi
A lactam is a cyclic amide. The term is a portmanteau of the words lactone + amide. Greek prefixes in alphabetical order indicate ring size: α-Lactam β-Lactam γ-Lactam δ-Lactam ε-Lactam This ring-size nomenclature stems from the fact that a hydrolyzed α-Lactam leads to an α-amino acid and a β-Lactam to a β-amino acid, etc. General synthetic methods exist for the organic synthesis of lactams. Lactams form by the acid-catalyzed rearrangement of oximes in the Beckmann rearrangement. Lactams form from hydrazoic acid in the Schmidt reaction. Lactams form from cyclisation of amino acids. Lactams form from intramolecular attack of linear acyl derivatives from the nucleophilic abstraction reaction. In iodolactamization an iminium ion reacts with a halonium ion formed in situ by reaction of an alkene with iodine. Lactams form by copper-catalyzed 1,3-dipolar cycloaddition of alkynes and nitrones in the Kinugasa reaction Diels-Alder reaction between cyclopentadiene and chlorosulfonyl isocyanate can be utilized to obtain both β- as well as γ-lactam.
At lower temp, β-lactam is the preferred product. At optimum temperatures, a useful γ-lactam known as Vince Lactam is obtained. A lactim is a cyclic carboximidic acid compound characterized by an endocyclic carbon-nitrogen double bond, they are formed. Lactams can polymerize to polyamides. Lactone, a cyclic ester. Β-Lactam β-Lactam antibiotics, which includes penicillins 2-Pyrrolidone 2-Piperidinone Caprolactam Media related to Lactams at Wikimedia Commons