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, hyperconjugation is the interaction of the electrons in a sigma orbital with an adjacent empty non-bonding or antibonding σ or π orbital to give an extended molecular orbital. Increased electron delocalization associated with hyperconjugation increases the stability of the system. Only electrons in bonds that are in the β position can have this sort of direct stabilizing effect—donating from a sigma bond on an atom to an orbital in another atom directly attached to it. However, extended versions of hyperconjugation can be important as well; the Baker–Nathan effect, sometimes used synonymously for hyperconjugation, is a specific application of it to certain chemical reactions or types of structures. Hyperconjugation can be used to rationalize a variety of other chemical phenomena, including the anomeric effect, the gauche effect, the rotational barrier of ethane, the beta-silicon effect, the vibrational frequency of exocyclic carbonyl groups, the relative stability of substituted carbocations and substituted carbon centred radicals.
Hyperconjugation is proposed by quantum mechanical modeling to be the correct explanation for the preference of the staggered conformation rather than the old textbook notion of steric hindrance. Hyperconjugation affects several properties. Bond length: Hyperconjugation is suggested as a key factor in shortening of sigma bonds. For example, the single C–C bonds in 1,3-butadiene and Propyne are 1.46 angstrom in length, much less than the value of around 1.54 Å found in saturated hydrocarbons. For butadiene, this can be explained as normal conjugation of the two alkenyl parts, but for Propyne, hyperconjugation between the alkyl and alkynyl parts. Dipole moments: The large increase in dipole moment of 1,1,1-trichloroethane as compared with chloroform can be attributed to hyperconjugated structures; the heat of formation of molecules with hyperconjugation are greater than sum of their bond energies and the heats of hydrogenation per double bond are less than the heat of hydrogenation of ethylene.
Stability of carbocations: 3C+ > 2CH+ > CH2+ > CH3+ The three C–H σ bonds of the methyl group attached to the carbocation can undergo the stabilization interaction but only one of them can be aligned with the empty p-orbital, depending on the conformation of the carbon–carbon bond. Donation from the two misaligned C–H bonds is weaker; the more adjacent methyl groups there are, the larger hyperconjugation stabilization is because of the increased number of adjacent C–H bonds. Early studies in hyperconjugation were performed by in the research group of George Kistiakowsky, their work, first published in 1937, was intended as a preliminary progress report of thermochemical studies of energy changes during addition reactions of various unsaturated and cyclic compounds. One set of experiments involved collected heats of hydrogenation data during gas-phase reactions of a range of compounds that contained one alkene unit; when comparing a range of monoalkyl-substituted alkenes, they found any alkyl group noticeably increased the stability, but that the choice of different specific alkyl groups had little to no effect.
A portion of Kistiakowsky’s work involved a comparison of other unsaturated compounds in the form of CH2=CHn-CH=CH2. These experiments revealed an important result; this is likened to the addition of two alkyl groups into ethylene. Kistiakowsky investigated open chain systems, where the largest value of heat liberated was found to be during the addition to a molecule in the 1,4-position. Cyclic molecules proved to be the most problematic, as it was found that the strain of the molecule would have to be considered; the strain of five-membered rings increased with a decrease degree of unsaturation. This was a surprising result, further investigated in work with cyclic acid anhydrides and lactones. Cyclic molecules like benzene and its derivatives were studied, as their behaviors were different from other unsaturated compounds. Despite the thoroughness of Kistiakowsky’s work, it was not complete and needed further evidence to back up his findings, his work was a crucial first step to the beginnings of the ideas of hyperconjugation and conjugation effects.
The conjugation of 1,3-butadiene was first evaluated by Kistiakowsky, a conjugative contribution of 3.5 kcal/mol was found based on the energetic comparison of hydrogenation between conjugated species and unconjugated analogues. Rogers who used the method first applied by Kistiakowsky, reported that the conjugation stabilization of 1,3-butadiyne was zero, as the difference of ΔhydH between first and second hydrogenation was zero; the heats of hydrogenation were obtained by computational G3 quantum chemistry method. Another group led by Houk suggested the methods employed by Rogers and Kistiakowsky was inappropriate, because that comparisons of heats of hydrogenation evaluate not only conjugation effects but other structural and electronic differences, they obtained -70.6 kcal/mol and -70.4 kcal/mol for the first and second hydrogenation by ab initio calculation, which confirmed Rogers’ data. However, they interpreted the data differently by taking into account the hyperconjugation stabilization.
To quantify hyperconjugation effect, they designed the following isodesmic reactions in 1-butyne and 1-butene. Deleting the hyperconjugative interactions gives virtual states that have energies that are 4.9 and 2.4 kcal/mol higher than those of 1-butyne and 1-butene, respectively. Employment of these virtual states results in a 9.6 kcal/mol conjugative stabilization for 1,3-butadiyne and 8.5 kcal/mol for
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
Three-center four-electron bond
The 3-center 4-electron bond is a model used to explain bonding in certain hypervalent molecules such as tetratomic and hexatomic interhalogen compounds, sulfur tetrafluoride, the xenon fluorides, the bifluoride ion. It is known as the Pimentel–Rundle three-center model after the work published by George C. Pimentel in 1951, which built on concepts developed earlier by Robert E. Rundle for electron-deficient bonding. An extended version of this model is used to describe the whole class of hypervalent molecules such as phosphorus pentafluoride and sulfur hexafluoride as well as multi-center π-bonding such as ozone and sulfur trioxide. While the term "hypervalent" was not introduced in the chemical literature until 1969, Irving Langmuir and G. N. Lewis debated the nature of bonding in hypervalent molecules as early as 1921. While Lewis supported the viewpoint of expanded octet, invoking s-p-d hybridized orbitals and maintaining 2c–2e bonds between neighboring atoms, Langmuir instead opted for maintaining the octet rule, invoking an ionic basis for bonding in hypervalent compounds.
In a 1951 seminal paper, Pimentel rationalized the bonding in hypervalent trihalide ions via a molecular orbital description, building on the concept of the "half-bond" introduced by Rundle in 1947. In this model, two of the four electrons occupy an all in-phase bonding MO, while the other two occupy a non-bonding MO, leading to an overall bond order of 0.5 between adjacent atoms. More recent theoretical studies on hypervalent molecules support the Langmuir view, confirming that the octet rule serves as a good first approximation to describing bonding in the s- and p-block elements. Triiodide Xenon difluoride Krypton difluoride Radon difluoride Argon fluorohydride Bifluoride Hydrogen Bonding Carboxylates Amides Ozone Azide Allyl anion The σ molecular orbitals of triiodide can be constructed by considering the in-phase and out-of-phase combinations of the central atom's p orbital with the p orbitals of the peripheral atoms; this exercise generates the diagram at right. Three molecular orbitals result from the combination of the three relevant atomic orbitals, with the four electrons occupying the two MOs lowest in energy – a bonding MO delocalized across all three centers, a non-bonding MO localized on the peripheral centers.
Using this model, one sidesteps the need to invoke hypervalent bonding considerations at the central atom, since the bonding orbital consists of two 2-center-1-electron bonds, the other two electrons occupy the non-bonding orbital. In the natural bond orbital viewpoint of 3c–4e bonding, the triiodide anion is constructed from the combination of the diiodine σ molecular orbitals and an iodide lone pair; the I− lone pair acts as a 2-electron donor, while the I2 σ* antibonding orbital acts as a 2-electron acceptor. Combining the donor and acceptor in in-phase and out-of-phase combinations results in the diagram depicted at right. Combining the donor lone pair with the acceptor σ* antibonding orbital results in an overall lowering in energy of the highest-occupied orbital. While the diagram depicted in Figure 2 shows the right-hand atom as the donor, an equivalent diagram can be constructed using the left-hand atom as the donor; this bonding scheme is succinctly summarized by the following two resonance structures: I—I···I− ↔ I−···I—I, which when averaged reproduces the I—I bond order of 0.5 obtained both from natural bond orbital analysis and from molecular orbital theory.
More recent theoretical investigations suggest the existence of a novel type of donor-acceptor interaction that may dominate in triatomic species with so-called "inverted electronegativity". Molecules of theoretical curiosity such as neon difluoride and berylium dilithide represent examples of inverted electronegativity; as a result of unusual bonding situation, the donor lone pair ends up with significant electron density on the central atom, while the acceptor is the "out-of-phase" combination of the p orbitals on the peripheral atoms. This bonding scheme is depicted in Figure 3 for the theoretical noble gas dihalide NeF2; the valence bond description and accompanying resonance structures A—B···C− ↔ A−···B—C suggest that molecules exhibiting 3c–4e bonding can serve as models for studying the transition states of bimolecular nucleophilic substitution reactions. Hypervalent molecule
A molecule is an electrically neutral group of two or more atoms held together by chemical bonds. Molecules are distinguished from ions by their lack of electrical charge. However, in quantum physics, organic chemistry, biochemistry, the term molecule is used less also being applied to polyatomic ions. In the kinetic theory of gases, the term molecule is used for any gaseous particle regardless of its composition. According to this definition, noble gas atoms are considered molecules as they are monatomic molecules. A molecule may be homonuclear, that is, it consists of atoms of one chemical element, as with oxygen. Atoms and complexes connected by non-covalent interactions, such as hydrogen bonds or ionic bonds, are not considered single molecules. Molecules as components of matter are common in organic substances, they make up most of the oceans and atmosphere. However, the majority of familiar solid substances on Earth, including most of the minerals that make up the crust and core of the Earth, contain many chemical bonds, but are not made of identifiable molecules.
No typical molecule can be defined for ionic crystals and covalent crystals, although these are composed of repeating unit cells that extend either in a plane or three-dimensionally. The theme of repeated unit-cellular-structure holds for most condensed phases with metallic bonding, which means that solid metals are not made of molecules. In glasses, atoms may be held together by chemical bonds with no presence of any definable molecule, nor any of the regularity of repeating units that characterizes crystals; the science of molecules is called molecular chemistry or molecular physics, depending on whether the focus is on chemistry or physics. Molecular chemistry deals with the laws governing the interaction between molecules that results in the formation and breakage of chemical bonds, while molecular physics deals with the laws governing their structure and properties. In practice, this distinction is vague. In molecular sciences, a molecule consists of a stable system composed of two or more atoms.
Polyatomic ions may sometimes be usefully thought of as electrically charged molecules. The term unstable molecule is used for reactive species, i.e. short-lived assemblies of electrons and nuclei, such as radicals, molecular ions, Rydberg molecules, transition states, van der Waals complexes, or systems of colliding atoms as in Bose–Einstein condensate. According to Merriam-Webster and the Online Etymology Dictionary, the word "molecule" derives from the Latin "moles" or small unit of mass. Molecule – "extremely minute particle", from French molécule, from New Latin molecula, diminutive of Latin moles "mass, barrier". A vague meaning at first; the definition of the molecule has evolved. Earlier definitions were less precise, defining molecules as the smallest particles of pure chemical substances that still retain their composition and chemical properties; this definition breaks down since many substances in ordinary experience, such as rocks and metals, are composed of large crystalline networks of chemically bonded atoms or ions, but are not made of discrete molecules.
Molecules are held together by ionic bonding. Several types of non-metal elements exist only as molecules in the environment. For example, hydrogen only exists as hydrogen molecule. A molecule of a compound is made out of two or more elements. A covalent bond is a chemical bond; these electron pairs are termed shared pairs or bonding pairs, the stable balance of attractive and repulsive forces between atoms, when they share electrons, is termed covalent bonding. Ionic bonding is a type of chemical bond that involves the electrostatic attraction between oppositely charged ions, is the primary interaction occurring in ionic compounds; the ions are atoms that have lost one or more electrons and atoms that have gained one or more electrons. This transfer of electrons is termed electrovalence in contrast to covalence. In the simplest case, the cation is a metal atom and the anion is a nonmetal atom, but these ions can be of a more complicated nature, e.g. molecular ions like NH4+ or SO42−. An ionic bond is the transfer of electrons from a metal to a non-metal for both atoms to obtain a full valence shell.
Most molecules are far too small to be seen with the naked eye. DNA, a macromolecule, can reach macroscopic sizes, as can molecules of many polymers. Molecules used as building blocks for organic synthesis have a dimension of a few angstroms to several dozen Å, or around one billionth of a meter. Single molecules cannot be observed by light, but small molecules and the outlines of individual atoms may be traced in some circumstances by use of an atomic force microscope; some of the largest molecules are supermolecules. The smallest molecule is the diatomic hydrogen, with a bond length of 0.74 Å. Effective molecular radius is the size; the table of permselectivity for different substances contains examples. The chemical formula for a molecule uses one line of chemical element symbols and sometimes al
Titanium disulfide is an inorganic compound with the formula TiS2. A golden yellow solid with high electrical conductivity, it belongs to a group of compounds called transition metal dichalcogenides, which consist of the stoichiometry ME2. TiS2 has been employed as a cathode material in rechargeable batteries. TiS2 adopts a hexagonal close packed structure, analogous to cadmium iodide. In this motif, half of the octahedral holes are filled with a "cation", in this case Ti4+; each Ti centre is surrounded by six sulfide ligands in an octahedral structure. Each sulfide is connected to the geometry at S being pyramidal. Several metal dichalcogenides adopt similar structures; the layers of TiS2 consist of covalent Ti-S bonds. The individual layers of TiS2 are bound together by van der Waals forces, which are weak intermolecular forces, it crystallises in the space group P3m1. The Ti-S bond lengths are 2.423 Å. The single most useful and most studied property of TiS2 is its ability to undergo intercalation upon treatment with electropositive elements.
The process is a redox reaction, illustrated in the case of lithium: TiS2 + Li → LiTiS2LiTiS2 is described as Li+. During the intercalation and deintercalation, a range of stiochimetries are produced with the general formul LixTiS2. During intercalation, the interlayer spacing expands and the electrical conductivity of the material increases. Intercalation is facilitated because of the weakness of the interlayer forces as well as the susceptibility of the Ti centers toward reduction. Intercalation can be conducted by combining a suspension of the disulfide material and a solution of the alkali metal in anhydrous ammonia. Alternatively solid TiS2 reacts with the alkali metal upon heating; the Rigid-Band Model, which assumes that electronic band structure does not change with intercalation, describes changes in the electronic properties upon intercalation. Deintercalation is the opposite of intercalation; this process is associated with recharging a Li/TiS2 battery. Intercalation and deintercalation can be monitored by cyclic voltammetry.
The microstructure of the titanium disulfide affects the intercalation and deintercalation kinetics. Titanium disulfide nanotubes have a higher uptake and discharge capacity than the polycrystalline structure; the higher surface area of the nanotubes is postulated to provide more binding sites for the anode ions than the polycrystalline structure. Formally containing the d0 ion Ti4+ and closed shell dianion S2−, TiS2 is diamagnetic, its magnetic susceptibility is the value being sensitive to stoichiometry. Titanium disulfide is a semimetal, meaning there is small overlap of the conduction band and valence band; the properties of titanium disulfide powder have been studied by high pressure synchrotron x-ray diffraction at room temperature. At ambient pressure, TiS2 behaves as semiconductor while at high pressures of 8 GPa the material behaves as a semimetal. At 15 GPa, the transport properties change. There is no significant change in the density of states at the Fermi level up to 20 GPa and phase change does not occur until 20.7 GPa.
A change in the structure of TiS2 was observed at a pressure of 26.3 GPa, although the new structure of the high pressure phase has not been determined. The unit cell of titanium disulfide is 3.407 by 5.695 angstroms. The size of the unit cell decreased at 17.8 GPa. The decrease in unit cell size was greater than was observed for MoS2 and WS2, indicating that titanium disulfide is softer and more compressible; the compression behavior of titanium disulfide is anisotropic. The axis parallel to S-Ti-S layers is more compressible than the axis perpendicular to S-Ti-S layers because of weak van der waals forces keeping S and Ti atoms together. At 17.8 GPa, the c-axis is compressed by 9.5% and the a-axis is compressed by 4%. The longitudinal sound velocity is 5284 m/s in the plane parallel to S-Ti-S layers; the longitudinal sound velocity perpendicular to the layers is 4383 m/s. Titanium disulfide is prepared by the reaction of the elements around 500 °C. Ti + 2 S → TiS2It can be more synthesized from titanium tetrachloride, but this product is less pure than that obtained from the elements.
TiCl4 + 2 H2S → TiS2 + 4 HClThis route has been applied to the formation of TiS2 films by chemical vapor deposition. Thiols and organic disulfides can be employed in place of hydrogen sulfide. Samples of TiS2 are unstable in air. Upon heating, the solid undergoes oxidation to titanium dioxide: TiS2 + O2 → TiO2 + 2STiS2 is sensitive to water: TiS2 + 2H2O → TiO2 + 2 H2SUpon heating, TiS2 releases sulfur, forming the titanium derivative: 2 TiS2 → Ti2S3 + S Thin films of TiS2 have been prepared by the sol-gel process from titanium isopropoxide followed by spin coating; this method affords amorphous material that crystallised at high temperatures to hexagonal TiS2, which crystallization orientations in the, directions. Because of their high surface area, such films are attractive for battery applications. More specialized morphologies - nanotubes, whiskers, thin films, fullerenes - are prepared by combining the standard reagents TiCl4 in unusual ways. For example, flower-like morphologies were obtain by treating a solution of sulfur in 1-octadecene with titanium tetrachloride.
A form of TiS2 with a fullerene-like structure has been prepared using the TiCl4/H2S method. The resulting spherical structures have diameters between 80 nm. Owing to their spherical shape, these fullerenes exhibit reduced friction c
A cathode is the electrode from which a conventional current leaves a polarized electrical device. This definition can be recalled by using the mnemonic CCD for Cathode Current Departs. A conventional current describes the direction. Electrons have a negative electrical charge, so the movement of electrons is opposite to that of the conventional current flow; the mnemonic cathode current departs means that electrons flow into the device's cathode from the external circuit. The electrode through which conventional current flows the other way, into the device, is termed an anode. Conventional current flow is from cathode to anode outside of the cell or device, regardless of the cell or device type and operating mode. Cathode polarity with respect to the anode can be positive or negative depending on how the device is being operated. Although positively charged cations always move towards the cathode and negatively charged anions move away from it, cathode polarity depends on the device type, can vary according to the operating mode.
In a device which absorbs energy of charge, the cathode is negative, in a device which provides energy, the cathode is positive: A battery or galvanic cell in use has a cathode, the positive terminal since, where the current flows out of the device. This outward current is carried internally by positive ions moving from the electrolyte to the positive cathode, it is continued externally by electrons moving into the battery which constitutes positive current flowing outwards. For example, the Daniell galvanic cell's copper electrode is the cathode. A battery, recharging or an electrolytic cell performing electrolysis has its cathode as the negative terminal, from which current exits the device and returns to the external generator as charge enters the battery/ cell. For example, reversing the current direction in a Daniell galvanic cell converts it into an electrolytic cell where the copper electrode is the positive terminal and the anode. In a diode, the cathode is the negative terminal at the pointed end of the arrow symbol, where current flows out of the device.
Note: electrode naming for diodes is always based on the direction of the forward current for types such as Zener diodes or solar cells where the current of interest is the reverse current. In vacuum tubes it is the negative terminal where electrons enter the device from the external circuit and proceed into the tube's near-vacuum, constituting a positive current flowing out of the device; the word was coined in 1834 from the Greek κάθοδος,'descent' or'way down', by William Whewell, consulted by Michael Faraday over some new names needed to complete a paper on the discovered process of electrolysis. In that paper Faraday explained that when an electrolytic cell is oriented so that electric current traverses the "decomposing body" in a direction "from East to West, or, which will strengthen this help to the memory, that in which the sun appears to move", the cathode is where the current leaves the electrolyte, on the West side: "kata downwards, `odos a way; the use of'West' to mean the'out' direction may appear unnecessarily contrived.
As related in the first reference cited above, Faraday had used the more straightforward term "exode". His motivation for changing it to something meaning'the West electrode' was to make it immune to a possible change in the direction convention for current, whose exact nature was not known at the time; the reference he used to this effect was the Earth's magnetic field direction, which at that time was believed to be invariant. He fundamentally defined his arbitrary orientation for the cell as being that in which the internal current would run parallel to and in the same direction as a hypothetical magnetizing current loop around the local line of latitude which would induce a magnetic dipole field oriented like the Earth's; this made the internal current East to West as mentioned, but in the event of a convention change it would have become West to East, so that the West electrode would not have been the'way out' any more. Therefore, "exode" would have become inappropriate, whereas "cathode" meaning'West electrode' would have remained correct with respect to the unchanged direction of the actual phenomenon underlying the current unknown but, he thought, unambiguously defined by the magnetic reference.
In retrospect the name change was unfortunate, not only because the Greek roots alone do not reveal the cathode's function any more, but more because, as we now know, the Earth's magnetic field direction on which the "cathode" term is based is subject to reversals whereas the current direction convention on which the "exode" term was based has no reason to change in the future. Since the discovery of the electron, an easier to remember, more durably technically correct, etymology has been suggested: cathode, from the Greek kathodos,'way down','the way into the cell for electrons'. In chemistry, a cathode is the electrode of an electrochemical cell.