Sulfuric acid
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
Nitroso
Nitroso refers to a functional group in organic chemistry which has the NO group attached to an organic moiety. As such, various nitroso groups can be categorized as C-nitroso compounds, S-nitroso compounds, N-nitroso compounds, O-nitroso compounds. Nitrosobenzene ( is prepared by reduction of nitrobenzene to phenylhydroxylamine, oxidized. Nitrosoarenes participate in a monomer-dimer equilibrium; the dimers, which are pale-yellow, are favored in the solid state, whereas the deep-green monomers are favored in dilute solution or at higher temperatures. They exist as cis- and trans-isomers. Nitroso compounds can be prepared by the reduction of nitro compounds or by the oxidation of hydroxylamines. A good example is 3CNO, known formally as 2-methyl-2-nitrosopropane, or t-BuNO. 3CNO is blue and exists in solution in equilibrium with its dimer, colorless, m.p. 80–81 °C. In the Fischer–Hepp rearrangement aromatic 4-nitrosoanilines are prepared from the corresponding nitrosamines. Another named reaction involving a nitroso compound is the Barton reaction.
Organonitroso compounds serve as a ligands for transition metals. Due to the stability of the nitric oxide free radical, nitroso organyls tend to have low C–N bond dissociation energies: nitrosoalkanes have BDEs on the order of 30 to 40 kcal/mol, while nitrosoarenes have BDEs on the order of 50 to 60 kcal/mol; as a consequence, they are heat- and light-sensitive. Compounds containing O– or N– bonds have lower bond dissociation energies. For instance, N-nitrosodiphenylamine, Ph2N–N=O, has a N–N bond dissociation energy of only 23 kcal/mol. Nitrite can enter two kinds of reaction, depending on the physico-chemical environment. Nitrosylation is adding a nitrosyl ion NO− to a metal or a thiol, leading to nitrosyl iron Fe–NO or S-nitrosothiols. Nitrosation is adding a nitrosonium ion NO+ to an amine –NH2 leading to a nitrosamine; this conversion occurs at acidic pH in the stomach, as shown in the equation for the formation of N-phenylnitrosamine: NO−2 + H+ ⇌ HONO HONO + H+ ⇌ H2O + NO+ C6H5NH2 + NO+ → C6H5NNO + H+Many primary alkyl N-nitroso compounds, such as CH3NNO, tend to be unstable with respect to hydrolysis to the alcohol.
Those derived from secondary amines are more robust. It is these N-nitrosamines. Nitrosyls are non-organic compounds containing the NO group, for example directly bound to the metal via the N atom, giving a metal–NO moiety. Alternatively, a nonmetal example is the common reagent nitrosyl chloride. Nitric oxide is a stable radical, having an unpaired electron. Reduction of nitric oxide gives the hyponitrite anion, NO−: NO + e− → NO−Oxidation of NO yields the nitrosonium cation, NO+: NO → NO+ + e−Nitric oxide can serve as a ligand forming metal nitrosyl complexes or just metal nitrosyls; these complexes can be viewed as adducts of NO −, or some intermediate case. In foodstuffs and in the gastro-intestinal tract and nitrosylation do not have the same consequences on consumer health. In cured meat: Meat processed by curing contains nitrite and has a pH of 5 where all nitrite is present as NO−2. Cured meat is added with sodium ascorbate; as demonstrated by S. Mirvish, ascorbate inhibits nitrosation of amines to nitrosamine, because ascorbate reacts with NO−2 to form NO.
Ascorbate and pH 5 thus favor nitrosylation of heme iron, forming nitrosylheme, a red pigment when included inside myoglobin, a pink pigment when it has been released by cooking. It participates to the "bacon flavor" of cured meat: nitrosylheme is thus considered a benefit for the meat industry and for consumers. In the stomach: secreted hydrogen chloride makes an acidic environment and ingested nitrite leads to nitrosation of amines, that yields nitrosamines. Nitrosation is low if vitamin C concentration is high. S-nitrosothiols are formed, that are stable at pH 2. In the colon: neutral pH does not favor nitrosation. No nitrosamine is formed in stools after addition of a secondary amine or nitrite. Neutral pH favors NO− release from S-nitrosothiols, nitrosylation of iron; the called NOC measured by Bingham's team in stools from red meat-fed volunteers were, according to Bingham and Kuhnle non-N-nitroso ATNC, e.g. S-nitrosothiols and nitrosyl iron. Nitrosamine, the functional group with the NO attached to an amine, such as R2N–NO Nitrosobenzene Nitric oxide Nitroxyl
Chlorine
Chlorine is a chemical element with symbol Cl and atomic number 17. The second-lightest of the halogens, it appears between fluorine and bromine in the periodic table and its properties are intermediate between them. Chlorine is a yellow-green gas at room temperature, it is an reactive element and a strong oxidising agent: among the elements, it has the highest electron affinity and the third-highest electronegativity on the Pauling scale, behind only oxygen and fluorine. The most common compound of chlorine, sodium chloride, has been known since ancient times. Around 1630, chlorine gas was first synthesised in a chemical reaction, but not recognised as a fundamentally important substance. Carl Wilhelm Scheele wrote a description of chlorine gas in 1774, supposing it to be an oxide of a new element. In 1809, chemists suggested that the gas might be a pure element, this was confirmed by Sir Humphry Davy in 1810, who named it from Ancient Greek: χλωρός, translit. Khlôros, lit.'pale green' based on its colour.
Because of its great reactivity, all chlorine in the Earth's crust is in the form of ionic chloride compounds, which includes table salt. It is the second-most abundant halogen and twenty-first most abundant chemical element in Earth's crust; these crustal deposits are dwarfed by the huge reserves of chloride in seawater. Elemental chlorine is commercially produced from brine by electrolysis; the high oxidising potential of elemental chlorine led to the development of commercial bleaches and disinfectants, a reagent for many processes in the chemical industry. Chlorine is used in the manufacture of a wide range of consumer products, about two-thirds of them organic chemicals such as polyvinyl chloride, many intermediates for the production of plastics and other end products which do not contain the element; as a common disinfectant, elemental chlorine and chlorine-generating compounds are used more directly in swimming pools to keep them clean and sanitary. Elemental chlorine at high concentrations is dangerous and poisonous for all living organisms, was used in World War I as the first gaseous chemical warfare agent.
In the form of chloride ions, chlorine is necessary to all known species of life. Other types of chlorine compounds are rare in living organisms, artificially produced chlorinated organics range from inert to toxic. In the upper atmosphere, chlorine-containing organic molecules such as chlorofluorocarbons have been implicated in ozone depletion. Small quantities of elemental chlorine are generated by oxidation of chloride to hypochlorite in neutrophils as part of the immune response against bacteria; the most common compound of chlorine, sodium chloride, has been known since ancient times. Its importance in food was well known in classical antiquity and was sometimes used as payment for services for Roman generals and military tribunes. Elemental chlorine was first isolated around 1200 with the discovery of aqua regia and its ability to dissolve gold, since chlorine gas is one of the products of this reaction: it was however not recognised as a new substance. Around 1630, chlorine was recognized as a gas by the Flemish chemist and physician Jan Baptist van Helmont.
The element was first studied in detail in 1774 by Swedish chemist Carl Wilhelm Scheele, he is credited with the discovery. Scheele produced chlorine by reacting MnO2 with HCl: 4 HCl + MnO2 → MnCl2 + 2 H2O + Cl2Scheele observed several of the properties of chlorine: the bleaching effect on litmus, the deadly effect on insects, the yellow-green color, the smell similar to aqua regia, he called it "dephlogisticated muriatic acid air" since it is a gas and it came from hydrochloric acid. He failed to establish chlorine as an element. Common chemical theory at that time held that an acid is a compound that contains oxygen, so a number of chemists, including Claude Berthollet, suggested that Scheele's dephlogisticated muriatic acid air must be a combination of oxygen and the yet undiscovered element, muriaticum. In 1809, Joseph Louis Gay-Lussac and Louis-Jacques Thénard tried to decompose dephlogisticated muriatic acid air by reacting it with charcoal to release the free element muriaticum, they did not succeed and published a report in which they considered the possibility that dephlogisticated muriatic acid air is an element, but were not convinced.
In 1810, Sir Humphry Davy tried the same experiment again, concluded that the substance was an element, not a compound. He announced his results to the Royal Society on 15 November that year. At that time, he named this new element "chlorine", from the Greek word χλωρος, meaning green-yellow; the name "halogen", meaning "salt producer", was used for chlorine in 1811 by Johann Salomo Christoph Schweigger. This term was used as a generic term to describe all the elements in the chlorine family, after a suggestion by Jöns Jakob Berzelius in 1826. In 1823, Michael Faraday liquefied chlorine for the first time, demonstrated that what was known as "solid chlorine" had a structure of chlorine hydrate. Chlorine gas was first used by French chemist Claude Berthollet to bleach textiles in 1785. Modern bleaches resulted from further work by Berthollet, who first produced sodium hypochlorite in 1789 in his laboratory in the town of Javel, by passing chlorine gas through a solution of sodium carbonate; the resulting liqu
Hydroxy group
A hydroxy or hydroxyl group is the entity with the formula OH. It contains oxygen bonded to hydrogen. In organic chemistry and carboxylic acids contain hydroxy groups; the anion, called hydroxide, consists of a hydroxyl group. According to IUPAC rules, the term hydroxyl refers to the radical OH only, while the functional group −OH is called hydroxy group. Water, carboxylic acids, many other hydroxy-containing compounds can be deprotonated readily; this behavior is rationalized by the disparate electronegativities of hydrogen. Hydroxy-containing compounds engage in hydrogen bonding, which causes them to stick together, leading to higher boiling and melting points than found for compounds that lack this functional group. Organic compounds, which are poorly soluble in water, become water-soluble when they contain two or more hydroxy groups, as illustrated by sugars and amino acid; the hydroxy group is pervasive in biochemistry. Many inorganic compounds contain hydroxy groups, including sulfuric acid, the chemical compound produced on the largest scale industrially.
Hydroxy groups participate in the dehydration reactions that link simple biological molecules into long chains. The joining of a fatty acid to glycerol to form a triacylglycerol removes the −OH from the carboxy end of the fatty acid; the joining of two aldehyde sugars to form a disaccharide removes the −OH from the carboxy group at the aldehyde end of one sugar. The creation of a peptide bond to link two amino acids to make a protein removes the −OH from the carboxy group of one amino acid. Hydroxyl radicals are reactive and undergo chemical reactions that make them short-lived; when biological systems are exposed to hydroxyl radicals, they can cause damage to cells, including those in humans, where they can react with DNA, proteins. In 2009, India's Chandrayaan-1 satellite, NASA's Cassini spacecraft and the Deep Impact probe have each detected the presence of water by evidence of hydroxyl fragments on the Moon; as reported by Richard Kerr, "A spectrometer detected an infrared absorption at a wavelength of 3.0 micrometers that only water or hydroxyl—a hydrogen and an oxygen bound together—could have created."
NASA reported in 2009 that the LCROSS probe revealed an ultraviolet emission spectrum consistent with hydroxyl presence. The Venus Express orbiter sent back Venus science data from April 2006 until December 2014. Results from Venus Express include the detection of hydroxyl in the atmosphere. Hydronium Ion Oxide Reece, Jane. "Unit 1, Chapter 4 &5." In Campbell Biology. Berge, Susan. San Francisco: Pearson Benjamin Cummings. ISBN 978-0-321-55823-7
Sulfonic acid
A sulfonic acid refers to a member of the class of organosulfur compounds with the general formula R−S2−OH, where R is an organic alkyl or aryl group and the S2 group a sulfonyl hydroxide. As a substituent, it is known as a sulfo group. A sulfonic acid can be thought of as sulfuric acid with one hydroxyl group replaced by an organic substituent; the parent compound is the parent sulfonic acid, HS2, a tautomer of sulfurous acid, S2. Salts or esters of sulfonic acids are called sulfonates. A sulfonic acid is produced by the process of sulfonation; the sulfonating agent is sulfur trioxide. A large scale application of this method is the production of alkylbenzenesulfonic acids: RC6H5 + SO3 → RC6H4SO3HIn this reaction, sulfur trioxide is an electrophile and the arene undergoes electrophilic aromatic substitution. Thiols can be oxidized to sulfonic acids: RSH + 3⁄2 O2 → RSO3HCertain sulfonic acids, such as perfluorooctanesulfonic acid, are prepared by electrophilic fluorination of preformed sulfonic acids.
The net conversion can be represented simplistically: C8H17SO3H + 17 F2 → C8F17SO3H + 17 HF Sulfonic acids are strong acids. They are cited as being around a million times stronger than the corresponding carboxylic acid. For example, p-Toluenesulfonic acid and methanesulfonic acid have pKa values of −2.8 and −1.9 while those of benzoic acid and acetic acid are 4.20 and 4.76, respectively. However, as a consequence of their strong acidity, their pKa values cannot be measured directly, values quoted should be regarded as indirect estimates with significant uncertainties. For instance, various sources have reported the pKa of methanesulfonic acid to be as high as −0.6 or as low as −6.5. Sulfonic acids are known to react with solid sodium chloride to form the sodium sulfonate and hydrogen chloride; this property implies an acidity within two or three orders of magnitude of that of HCl, whose pKa was accurately determined. Because of their polarity, sulfonic acids tend to be crystalline solids or viscous, high-boiling liquids.
They are usually colourless and nonoxidizing, which makes them suitable for use as acid catalysts in organic reactions. Their polarity, in conjunction with their high acidity, renders short-chain sulfonic acids water soluble, while longer-chain ones exhibit detergent-like properties; the structure of sulfonic acids is illustrated by methanesulfonic acid. The sulfonic acid group, RSO2OH features a tetrahedral sulfur centre, meaning that sulfur is at the center of four atoms: three oxygens and one carbon; the overall geometry of the sulfur centre is reminiscent of the shape of sulfuric acid. Representative sulfonic acids and sulfonates Although both alkyl and aryl sulfonic acids are known, most of the applications are associated with the aromatic derivatives. Detergents and surfactants are molecules that combine nonpolar and polar groups. Traditionally, soaps are the popular surfactants, being derived from fatty acids. Since the mid-20th century, the usage of sulfonic acids has surpassed soap in advanced societies.
For example, an estimated 2 billion kilograms of alkylbenzenesulfonates are produced annually for diverse purposes. Lignin sulfonates, produced by sulfonation of lignin are components of drilling fluids and additives in certain kinds of concrete. Many if not most of the anthroquinone dyes are processed via sulfonation. Sulfonic acids tend to bind to proteins and carbohydrates. Most "washable" dyes are sulfonic acids for this reason. P-Cresidinesulfonic acid is used to make food dyes. Being strong acids, sulfonic acids are used as catalysts; the simplest examples are methanesulfonic acid, CH3SO2OH and p-toluenesulfonic acid, which are used in organic chemistry as acids that are lipophilic. Polymeric sulfonic acids are useful. Dowex resin are sulfonic acid derivatives of polystyrene and is used as catalysts and for ion exchange. Nafion, a fluorinated polymeric sulfonic acid is a component of proton exchange membranes in fuel cells. Sulfa drugs, a class of antibacterials, are produced from sulfonic acids.
Methanesulfonic acid is used as the supporting electrolyte of the zinc-cerium and lead-acid flow batteries. Arylsulfonic acids are susceptible to the reverse of the sulfonation reaction. Whereas benzene sulfonic acid hydrolyzes above 200 ″C, most related derivatives are easier to hydrolyze. Thus, heating aryl sulfonic acids in aqueous acid produces the parent arene; this reaction is employed in several scenarios. In some cases the sulfonic acid serves as a water-solubilizing protecting group, as illustrated by the purification of para-xylene via its sulfonic acid derivative; the synthesis of 2,6-dichlorophenol, phenol is converted to its 4-sulfonic acid derivative, which selectively chlorinates at the positions flanking the phenol. Hydrolysis releases the sulfonic acid group. Sulfonic acids can be converted to esters; this class of organic compounds has the general formula R−SO2−OR. Sulfonic esters such as methyl triflate are considered good alkylating agents in organic synthesis; such sulfonate esters are prepared by alcoholysis of the sulfonyl chlorides: RSO2Cl + R′OH → RSO2OR′ + HCl Sulfonyl halide groups occur when a sulfonyl functional group is singly bonded to a halogen atom.
They have the general formula R−SO2−X where X is a halide invariably chloride. They are produced by chlorination of sulfonic acids using related reagents. Although strong, the C−SO3- bond can be broken by nucleophilic reagents. Of historic and
Scanning tunneling microscope
A scanning tunneling microscope is an instrument for imaging surfaces at the atomic level. Its development in 1981 earned its inventors, Gerd Binnig and Heinrich Rohrer, the Nobel Prize in Physics in 1986. For an STM, good resolution is considered to be 0.1 nm lateral resolution and 0.01 nm depth resolution. With this resolution, individual atoms within materials are imaged and manipulated; the STM can be used not only in ultra-high vacuum but in air and various other liquid or gas ambients, at temperatures ranging from near zero kelvin to over 1000 °C. STM is based on the concept of quantum tunneling; when a conducting tip is brought near to the surface to be examined, a bias applied between the two can allow electrons to tunnel through the vacuum between them. The resulting tunneling current is a function of tip position, applied voltage, the local density of states of the sample. Information is acquired by monitoring the current as the tip's position scans across the surface, is displayed in image form.
STM can be a challenging technique, as it requires clean and stable surfaces, sharp tips, excellent vibration control, sophisticated electronics, but nonetheless many hobbyists have built their own. First, a voltage bias is applied and the tip is brought close to the sample by coarse sample-to-tip control, turned off when the tip and sample are sufficiently close. At close range, fine control of the tip in all three dimensions when near the sample is piezoelectric, maintaining tip-sample separation W in the 4-7 Å range, the equilibrium position between attractive and repulsive interactions. In this situation, the voltage bias will cause electrons to tunnel between the tip and sample, creating a current that can be measured. Once tunneling is established, the tip's bias and position with respect to the sample can be varied and data are obtained from the resulting changes in current. If the tip is moved across the sample in the x-y plane, the changes in surface height and density of states causes changes in current.
These changes are mapped in images. This change in current with respect to position can be measured itself, or the height, z, of the tip corresponding to a constant current can be measured; these two modes are called constant current mode, respectively. In constant current mode, feedback electronics adjust the height by a voltage to the piezoelectric height control mechanism; this leads to a height variation and thus the image comes from the tip topography across the sample and gives a constant charge density surface. In constant height mode, the voltage and height are both held constant while the current changes to keep the voltage from changing; the benefit to using a constant height mode is that it is faster, as the piezoelectric movements require more time to register the height change in constant current mode than the current change in constant height mode. All images produced by STM are grayscale, with color optionally added in post-processing in order to visually emphasize important features.
In addition to scanning across the sample, information on the electronic structure at a given location in the sample can be obtained by sweeping voltage and measuring current at a specific location. This type of measurement is called scanning tunneling spectroscopy and results in a plot of the local density of states as a function of energy within the sample; the advantage of STM over other measurements of the density of states lies in its ability to make local measurements: for example, the density of states at an impurity site can be compared to the density of states far from impurities. Framerates of at least 25 Hz enable so called video-rate STM. Framerates up to 80 Hz are possible with working feedback that adjusts the height of the tip. Due to the line-by-line scanning motion, a proper comparison on the speed requires not only the framerate, but the number of pixels in an image: with a framerate of 10 Hz and 100x100 pixels the tip moves with a line frequency of 1 kHz, whereas it moves with only with 500 Hz, when measuring with a faster framerate of 50 Hz but only 10x10 pixels.
Video-rate STM can be used to scan surface diffusion. The components of an STM include scanning tip, piezoelectric controlled height and x,y scanner, coarse sample-to-tip control, vibration isolation system, computer; the resolution of an image is limited by the radius of curvature of the scanning tip of the STM. Additionally, image artifacts can occur if the tip has two tips at the end rather than a single atom. Therefore, it has been essential to develop processes for obtaining sharp, usable tips. Carbon nanotubes have been used in this instance; the tip is made of tungsten or platinum-iridium, though gold is used. Tungsten tips are made by electrochemical etching, platinum-iridium tips by mechanical shearing. Due to the extreme sensitivity of tunnel current to height, proper vibration insulation or an rigid STM body is imperative for obtaining usable results. In the first STM by Binnig and Rohrer, magnetic levitation was used to keep the STM free from vibrations. Additionally, mechanisms for reducing eddy currents are sometimes implemented.
Maintaining the tip position with respect to the sample, scanning the
Delocalized electron
In chemistry, delocalized electrons are electrons in a molecule, ion or solid metal that are not associated with a single atom or a covalent bond. The term is general and can have different meanings in different fields. In organic chemistry, this refers to resonance in aromatic compounds. In solid-state physics, this refers to free electrons. In quantum chemistry, this refers to molecular orbital electrons that extend over several adjacent atoms. In the simple aromatic ring of benzene the delocalization of six π electrons over the C6 ring is graphically indicated by a circle; the fact that the six C-C bonds are equidistant is one indication that the π electrons are delocalized. In valence bond theory, delocalization in benzene is represented by resonance structures. Delocalized electrons exist in the structure of solid metals. Metallic structure consists of aligned positive ions in a "sea" of delocalized electrons; this means that the electrons are free to move throughout the structure, gives rise to properties such as conductivity.
In diamond all four outer electrons of each carbon atom are'localized' between the atoms in covalent bonding. The movement of electrons is restricted and diamond does not conduct an electric current. In graphite, each carbon atom uses only 3 of its 4 outer energy level electrons in covalently bonding to three other carbon atoms in a plane; each carbon atom contributes one electron to a delocalized system of electrons, a part of the chemical bonding. The delocalized electrons are free to move throughout the plane. For this reason, graphite conducts electricity along the planes of carbon atoms, but does not conduct in a direction at right angles to the plane. Standard ab initio quantum chemistry methods lead to delocalized orbitals that, in general, extend over an entire molecule and have the symmetry of the molecule. Localized orbitals may be found as linear combinations of the delocalized orbitals, given by an appropriate unitary transformation. In the methane molecule for example, ab initio calculations show bonding character in four molecular orbitals, sharing the electrons uniformly among all five atoms.
There are two orbital levels, a bonding molecular orbital formed from the 2s orbital on carbon and triply degenerate bonding molecular orbitals from each of the 2p orbitals on carbon. The localized sp3 orbitals corresponding to each individual bond in valence bond theory can be obtained from a linear combination of the four molecular orbitals. Aromatic ring current Electride Solvated electron
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