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
Hydrochloric acid or muriatic acid is a colorless inorganic chemical system with the formula H2O:HCl. Hydrochloric acid has a distinctive pungent smell, it is classified as acidic and can attack the skin over a wide composition range, since the hydrogen chloride dissociates in aqueous solution. Hydrochloric acid is the simplest chlorine-based acid system containing water, it is a solution of hydrogen chloride and water, a variety of other chemical species, including hydronium and chloride ions. It is an important chemical reagent and industrial chemical, used in the production of polyvinyl chloride for plastic. In households, diluted hydrochloric acid is used as a descaling agent. In the food industry, hydrochloric acid is used in the production of gelatin. Hydrochloric acid is used in leather processing. Hydrochloric acid was discovered by the alchemist Jabir ibn Hayyan around the year 800 AD, it was called acidum salis and spirits of salt because it was produced from rock salt and "green vitriol" and from the chemically similar common salt and sulfuric acid.
Free hydrochloric acid was first formally described in the 16th century by Libavius. It was used by chemists such as Glauber and Davy in their scientific research. Unless pressurized or cooled, hydrochloric acid will turn into a gas if there is around 60% or less of water. Hydrochloric acid is known as hydronium chloride, in contrast to its anhydrous parent known as hydrogen chloride, or dry HCl. Hydrochloric acid was known to European alchemists as spirits of acidum salis. Both names are still used in other languages, such as German: Salzsäure, Dutch: Zoutzuur, Swedish: Saltsyra, Turkish: Tuz Ruhu, Polish: kwas solny, Bulgarian: солна киселина, Russian: соляная кислота, Chinese: 鹽酸, Korean: 염산, Taiwanese: iâm-sng. Gaseous HCl was called marine acid air; the old name muriatic acid has the same origin, this name is still sometimes used. The name hydrochloric acid was coined by the French chemist Joseph Louis Gay-Lussac in 1814. Hydrochloric acid has been an important and used chemical from early history and was discovered by the alchemist Jabir ibn Hayyan around the year 800 AD.
Aqua regia, a mixture consisting of hydrochloric and nitric acids, prepared by dissolving sal ammoniac in nitric acid, was described in the works of Pseudo-Geber, a 13th-century European alchemist. Other references suggest that the first mention of aqua regia is in Byzantine manuscripts dating to the end of the 13th century. Free hydrochloric acid was first formally described in the 16th century by Libavius, who prepared it by heating salt in clay crucibles. Other authors claim that pure hydrochloric acid was first discovered by the German Benedictine monk Basil Valentine in the 15th century, when he heated common salt and green vitriol, whereas others argue that there is no clear reference to the preparation of pure hydrochloric acid until the end of the 16th century. In the 17th century, Johann Rudolf Glauber from Karlstadt am Main, Germany used sodium chloride salt and sulfuric acid for the preparation of sodium sulfate in the Mannheim process, releasing hydrogen chloride gas. Joseph Priestley of Leeds, England prepared pure hydrogen chloride in 1772, by 1808 Humphry Davy of Penzance, England had proved that the chemical composition included hydrogen and chlorine.
During the Industrial Revolution in Europe, demand for alkaline substances increased. A new industrial process developed by Nicolas Leblanc of Issoudun, France enabled cheap large-scale production of sodium carbonate. In this Leblanc process, common salt is converted to soda ash, using sulfuric acid and coal, releasing hydrogen chloride as a by-product; until the British Alkali Act 1863 and similar legislation in other countries, the excess HCl was vented into the air. After the passage of the act, soda ash producers were obliged to absorb the waste gas in water, producing hydrochloric acid on an industrial scale. In the 20th century, the Leblanc process was replaced by the Solvay process without a hydrochloric acid by-product. Since hydrochloric acid was fully settled as an important chemical in numerous applications, the commercial interest initiated other production methods, some of which are still used today. After the year 2000, hydrochloric acid is made by absorbing by-product hydrogen chloride from industrial organic compounds production.
Since 1988, hydrochloric acid has been listed as a Table II precursor under the 1988 United Nations Convention Against Illicit Traffic in Narcotic Drugs and Psychotropic Substances because of its use in the production of heroin and methamphetamine. Hydrochloric acid is the salt of H3O + and chloride, it is prepared by treating HCl with water. HCl + H 2 O ⟶ H 3 O + + Cl − However, the speciation of hydrochloric acid is more complicated than this simple equation implies; the structure of bulk water is infamously complex, the formula H3O+ is a gross oversimplification of the true nature of the solvated proton, H+, present in hydrochloric acid. A combined IR, Raman, X-ray and neutron diffraction study of concentrated solutions of hydrochloric acid revealed that the primary form of H+ in these solutions is H5O2+, along with the chloride anion, is hydrogen-bonded to neighboring wa
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
In organic chemistry, an alkyne is an unsaturated hydrocarbon containing at least one carbon—carbon triple bond. The simplest acyclic alkynes with only one triple bond and no other functional groups form a homologous series with the general chemical formula CnH2n−2. Alkynes are traditionally known as acetylenes, although the name acetylene refers to C2H2, known formally as ethyne using IUPAC nomenclature. Like other hydrocarbons, alkynes are hydrophobic but tend to be more reactive. Alkynes are characteristically more unsaturated than alkenes, thus they add two equivalents of bromine. Other reactions are listed below. In some reactions, alkynes are less reactive than alkenes. For example, in a molecule with an -ene and an -yne group, addition occurs preferentially at the -ene. Possible explanations involve the two π-bonds in the alkyne delocalising, which would reduce the energy of the π-system or the stability of the intermediates during the reaction, they show greater tendency to oligomerize than alkenes do.
The resulting polymers, called polyacetylenes are conjugated and can exhibit semiconducting properties. In acetylene, the H–C≡C bond angles are 180°. By virtue of this bond angle, alkynes are rod-like. Correspondingly, cyclic alkynes are rare. Benzyne is unstable; the C≡C bond distance of 121 picometers is much shorter than the C=C distance in alkenes or the C–C bond in alkanes. The triple bond is strong with a bond strength of 839 kJ/mol; the sigma bond contributes 369 kJ/mol, the first pi bond contributes 268 kJ/mol and the second pi-bond of 202 kJ/mol bond strength. Bonding discussed in the context of molecular orbital theory, which recognizes the triple bond as arising from overlap of s and p orbitals. In the language of valence bond theory, the carbon atoms in an alkyne bond are sp hybridized: they each have two unhybridized p orbitals and two sp hybrid orbitals. Overlap of an sp orbital from each atom forms one sp–sp sigma bond; each p orbital on one atom overlaps one on the other atom, forming two pi bonds, giving a total of three bonds.
The remaining sp orbital on each atom can form a sigma bond to another atom, for example to hydrogen atoms in the parent acetylene. The two sp orbitals project on opposite sides of the carbon atom. Internal alkynes feature carbon substituents on each acetylenic carbon. Symmetrical examples include 3-hexyne. Terminal alkynes have the formula RC2H. An example is methylacetylene. Terminal alkynes, like acetylene itself, are mildly acidic, with pKa values of around 25, they are far more acidic than alkenes and alkanes, which have pKa values of around 40 and 50, respectively. The acidic hydrogen on terminal alkynes can be replaced by a variety of groups resulting in halo-, silyl-, alkoxoalkynes; the carbanions generated by deprotonation of terminal alkynes are called acetylides. In systematic chemical nomenclature, alkynes are named with the Greek prefix system without any additional letters. Examples include octyne. In parent chains with four or more carbons, it is necessary to say. For octyne, one can either write oct-3-yne when the bond starts at the third carbon.
The lowest number possible is given to the triple bond. When no superior functional groups are present, the parent chain must include the triple bond if it is not the longest possible carbon chain in the molecule. Ethyne is called by its trivial name acetylene. In chemistry, the suffix -yne is used to denote the presence of a triple bond. In organic chemistry, the suffix follows IUPAC nomenclature. However, inorganic compounds featuring unsaturation in the form of triple bonds may be denoted by substitutive nomenclature with the same methods used with alkynes. "-diyne" is used when there are two triple bonds, so on. The position of unsaturation is indicated by a numerical locant preceding the "-yne" suffix, or'locants' in the case of multiple triple bonds. Locants are chosen. "-yne" is used as an infix to name substituent groups that are triply bound to the parent compound. Sometimes a number between hyphens is inserted before it to state which atoms the triple bond is between; this suffix arose as a collapsed form of the end of the word "acetylene".
The final" - e" disappears. Commercially, the dominant alkyne is acetylene itself, used as a fuel and a precursor to other compounds, e.g. acrylates. Hundreds of millions of kilograms are produced annually by partial oxidation of natural gas: 2 CH4 + 3/2 O2 → HC≡CH + 3 H2OPropyne industrially useful, is prepared by thermal cracking of hydrocarbons. Most other industrially useful alkyne derivatives are prepared from acetylene, e.g. via condensation with formaldehyde. Specialty alkynes are prepared by dehydrohalogenation of vicinal alkyl dihalides or vinyl halides. Metal acetylides can be coupled with primary alkyl halides. Via the Fritsch–Buttenberg–Wiechell rearrangement, alkynes are prepared from vinyl bromides. Alkynes can be prepared from aldehydes using the Corey–Fuchs reaction and from aldehydes or ketones by the Seyferth–Gilbert homologation. In the alkyne zipper reaction, alkynes are generated from other alkynes by treatment with a strong base. Featuring a reactive functional group, alkynes participate in many organic reactions.
Such use was pioneered by Ralph Raphael, who in 1955 wrote the first book describing their versatility as intermediates in synthesis. Alkynes character
Bamberger triazine synthesis
The Bamberger triazine synthesis in organic chemistry is a classic organic synthesis of a triazine first reported by Eugen Bamberger in 1892. The reactants are an aryl diazonium salt obtained from reaction of the corresponding aniline with sodium nitrite and hydrochloric acid and the hydrazone of pyruvic acid; the azo intermediate converts to the benzotriazine in the third step with sulfuric acid in acetic acid. From the same inventor: the Bamberger rearrangement
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 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