Diamond is a solid form of the element carbon with its atoms arranged in a crystal structure called diamond cubic. At room temperature and pressure, another solid form of carbon known as graphite is the chemically stable form, but diamond never converts to it. Diamond has the highest hardness and thermal conductivity of any natural material, properties that are utilized in major industrial applications such as cutting and polishing tools, they are the reason that diamond anvil cells can subject materials to pressures found deep in the Earth. Because the arrangement of atoms in diamond is rigid, few types of impurity can contaminate it. Small numbers of defects or impurities color diamond blue, brown, purple, orange or red. Diamond has high optical dispersion. Most natural diamonds have ages between 1 billion and 3.5 billion years. Most were formed at depths between 150 and 250 kilometers in the Earth's mantle, although a few have come from as deep as 800 kilometers. Under high pressure and temperature, carbon-containing fluids dissolved minerals and replaced them with diamonds.
Much more they were carried to the surface in volcanic eruptions and deposited in igneous rocks known as kimberlites and lamproites. Synthetic diamonds can be grown from high-purity carbon under high pressures and temperatures or from hydrocarbon gas by chemical vapor deposition. Imitation diamonds can be made out of materials such as cubic zirconia and silicon carbide. Natural and imitation diamonds are most distinguished using optical techniques or thermal conductivity measurements. Diamond is a solid form of pure carbon with its atoms arranged in a crystal. Solid carbon comes in different forms known as allotropes depending on the type of chemical bond; the two most common allotropes of pure carbon are graphite. In graphite the bonds are sp2 orbital hybrids and the atoms form in planes with each bound to three nearest neighbors 120 degrees apart. In diamond they are sp3 and the atoms form tetrahedra with each bound to four nearest neighbors. Tetrahedra are rigid, the bonds are strong, of all known substances diamond has the greatest number of atoms per unit volume, why it is both the hardest and the least compressible.
It has a high density, ranging from 3150 to 3530 kilograms per cubic metre in natural diamonds and 3520 kg/m³ in pure diamond. In graphite, the bonds between nearest neighbors are stronger but the bonds between planes are weak, so the planes can slip past each other. Thus, graphite is much softer than diamond. However, the stronger bonds make graphite less flammable. Diamonds have been adapted for many uses because of the material's exceptional physical characteristics. Most notable are its extreme hardness and thermal conductivity, as well as wide bandgap and high optical dispersion. Diamond's ignition point is 720 -- 800 °C in 850 -- 1000 °C in air; the equilibrium pressure and temperature conditions for a transition between graphite and diamond is well established theoretically and experimentally. The pressure changes linearly between 1.7 GPa at 0 K and 12 GPa at 5000 K. However, the phases have a wide region about this line where they can coexist. At normal temperature and pressure, 20 °C and 1 standard atmosphere, the stable phase of carbon is graphite, but diamond is metastable and its rate of conversion to graphite is negligible.
However, at temperatures above about 4500 K, diamond converts to graphite. Rapid conversion of graphite to diamond requires pressures well above the equilibrium line: at 2000 K, a pressure of 35 GPa is needed. Above the triple point, the melting point of diamond increases with increasing pressure. At high pressures and germanium have a BC8 body-centered cubic crystal structure, a similar structure is predicted for carbon at high pressures. At 0 K, the transition is predicted to occur at 1100 GPa; the most common crystal structure of diamond is called diamond cubic. It is formed of unit cells stacked together. Although there are 18 atoms in the figure, each corner atom is shared by eight unit cells and each atom in the center of a face is shared by two, so there are a total of eight atoms per unit cell; each side of the unit cell is 3.57 angstroms in length. A diamond cubic lattice can be thought of as two interpenetrating face-centered cubic lattices with one displaced by 1/4 of the diagonal along a cubic cell, or as one lattice with two atoms associated with each lattice point.
Looked at from a <1 1 1> crystallographic direction, it is formed of layers stacked in a repeating ABCABC... pattern. Diamonds can form an ABAB... structure, known as hexagonal diamond or lonsdaleite, but this is far less common and is formed under different conditions from cubic carbon. Diamonds occur most as euhedral or rounded octahedra and twinned octahedra known as macles; as diamond's crystal structure has a cubic arrangement of the atoms, they have many facets that belong to a cube, rhombicosidodecahedron, tetrakis hexahedron or disdyakis dodecahedron. The crystals can be elongated. Diamonds are found coated in nyf, an opaque gum-like skin; some diamonds have opaque fibers. They are referred to as opaque if the fibers
Iron is a chemical element with symbol Fe and atomic number 26. It is a metal, that belongs to group 8 of the periodic table, it is by mass the most common element on Earth, forming much of Earth's inner core. It is the fourth most common element in the Earth's crust. Pure iron is rare on the Earth's crust being limited to meteorites. Iron ores are quite abundant, but extracting usable metal from them requires kilns or furnaces capable of reaching 1500 °C or higher, about 500 °C higher than what is enough to smelt copper. Humans started to dominate that process in Eurasia only about 2000 BCE, iron began to displace copper alloys for tools and weapons, in some regions, only around 1200 BCE; that event is considered the transition from the Bronze Age to the Iron Age. Iron alloys, such as steel and special steels are now by far the most common industrial metals, because of their mechanical properties and their low cost. Pristine and smooth pure iron surfaces are mirror-like silvery-gray. However, iron reacts with oxygen and water to give brown to black hydrated iron oxides known as rust.
Unlike the oxides of some other metals, that form passivating layers, rust occupies more volume than the metal and thus flakes off, exposing fresh surfaces for corrosion. The body of an adult human contains about 3 to 5 grams of elemental iron in hemoglobin and myoglobin; these two proteins play essential roles in vertebrate metabolism oxygen transport by blood and oxygen storage in muscles. To maintain the necessary levels, human iron metabolism requires a minimum of iron in the diet. Iron is the metal at the active site of many important redox enzymes dealing with cellular respiration and oxidation and reduction in plants and animals. Chemically, the most common oxidation states of iron are +2 and +3. Iron shares many properties of other transition metals, including the other group 8 elements and osmium. Iron forms compounds in a wide range of oxidation states, −2 to +7. Iron forms many coordination compounds. At least four allotropes of iron are known, conventionally denoted α, γ, δ, ε; the first three forms are observed at ordinary pressures.
As molten iron cools past its freezing point of 1538 °C, it crystallizes into its δ allotrope, which has a body-centered cubic crystal structure. As it cools further to 1394 °C, it changes to its γ-iron allotrope, a face-centered cubic crystal structure, or austenite. At 912 °C and below, the crystal structure again becomes the bcc α-iron allotrope; the physical properties of iron at high pressures and temperatures have been studied extensively, because of their relevance to theories about the cores of the Earth and other planets. Above 10 GPa and temperatures of a few hundred kelvin or less, α-iron changes into another hexagonal close-packed structure, known as ε-iron; the higher-temperature γ-phase changes into ε-iron, but does so at higher pressure. Some controversial experimental evidence exists for a stable β phase at pressures above 50 GPa and temperatures of at least 1500 K, it is supposed to have a double hcp structure. The inner core of the Earth is presumed to consist of an iron-nickel alloy with ε structure.
The melting and boiling points of iron, along with its enthalpy of atomization, are lower than those of the earlier 3d elements from scandium to chromium, showing the lessened contribution of the 3d electrons to metallic bonding as they are attracted more and more into the inert core by the nucleus. This same trend appears for ruthenium but not osmium; the melting point of iron is experimentally well defined for pressures less than 50 GPa. For greater pressures, published data still varies by tens of gigapascals and over a thousand kelvin. Below its Curie point of 770 °C, α-iron changes from paramagnetic to ferromagnetic: the spins of the two unpaired electrons in each atom align with the spins of its neighbors, creating an overall magnetic field; this happens because the orbitals of those two electrons do not point toward neighboring atoms in the lattice, therefore are not involved in metallic bonding. In the absence of an external source of magnetic field, the atoms get spontaneously partitioned into magnetic domains, about 10 micrometres across, such that the atoms in each domain have parallel spins, but different domains have other orientations.
Thus a macroscopic piece of iron will have a nearly zero overall magnetic field. Application of an external magnetic field causes the domains that are magnetized in the same general direction to grow at the expense of adjacent ones that point in other directions, reinforcing the external field; this effect is exploited in devices that needs to channel magnetic fields, such as electrical transformers, magnetic recording heads, electric motors. Impurities, lattice defects, or grain and particle boundaries can "pin" the domains in the new positions, so that the effect persists after the external field is removed -- thus turning the iron object into a magnet. Similar behavior is exhibited by some iron compounds, such as the fer
In geometry, a tetrahedron known as a triangular pyramid, is a polyhedron composed of four triangular faces, six straight edges, four vertex corners. The tetrahedron is the simplest of all the ordinary convex polyhedra and the only one that has fewer than 5 faces; the tetrahedron is the three-dimensional case of the more general concept of a Euclidean simplex, may thus be called a 3-simplex. The tetrahedron is one kind of pyramid, a polyhedron with a flat polygon base and triangular faces connecting the base to a common point. In the case of a tetrahedron the base is a triangle, so a tetrahedron is known as a "triangular pyramid". Like all convex polyhedra, a tetrahedron can be folded from a single sheet of paper, it has two such nets. For any tetrahedron there exists a sphere on which all four vertices lie, another sphere tangent to the tetrahedron's faces. A regular tetrahedron is one, it is one of the five regular Platonic solids. In a regular tetrahedron, all faces are the same size and shape and all edges are the same length.
Regular tetrahedra alone do not tessellate, but if alternated with regular octahedra in the ratio of two tetrahedra to one octahedron, they form the alternated cubic honeycomb, a tessellation. The regular tetrahedron is self-dual; the compound figure comprising two such dual tetrahedra form a stellated octahedron or stella octangula. The following Cartesian coordinates define the four vertices of a tetrahedron with edge length 2, centered at the origin, two level edges: and Expressed symmetrically as 4 points on the unit sphere, centroid at the origin, with lower face level, the vertices are: v1 = v2 = v3 = v4 = with the edge length of sqrt. Still another set of coordinates are based on an alternated cube or demicube with edge length 2; this form has Coxeter diagram and Schläfli symbol h. The tetrahedron in this case has edge length 2√2. Inverting these coordinates generates the dual tetrahedron, the pair together form the stellated octahedron, whose vertices are those of the original cube. Tetrahedron:, Dual tetrahedron:, For a regular tetrahedron of edge length a: With respect to the base plane the slope of a face is twice that of an edge, corresponding to the fact that the horizontal distance covered from the base to the apex along an edge is twice that along the median of a face.
In other words, if C is the centroid of the base, the distance from C to a vertex of the base is twice that from C to the midpoint of an edge of the base. This follows from the fact that the medians of a triangle intersect at its centroid, this point divides each of them in two segments, one of, twice as long as the other. For a regular tetrahedron with side length a, radius R of its circumscribing sphere, distances di from an arbitrary point in 3-space to its four vertices, we have d 1 4 + d 2 4 + d 3 4 + d 4 4 4 + 16 R 4 9 = 2.
Photochemistry is the branch of chemistry concerned with the chemical effects of light. This term is used to describe a chemical reaction caused by absorption of ultraviolet, visible light or infrared radiation. In nature, photochemistry is of immense importance as it is the basis of photosynthesis and the formation of vitamin D with sunlight. Photochemical reactions proceed differently than temperature-driven reactions. Photochemical paths access high energy intermediates that cannot be generated thermally, thereby overcoming large activation barriers in a short period of time, allowing reactions otherwise inaccessible by thermal processes. Photochemistry is destructive, as illustrated by the photodegradation of plastics. Photoexcitation is the first step in a photochemical process where the reactant is elevated to a state of higher energy, an excited state; the first law of photochemistry, known as the Grotthuss–Draper law, states that light must be absorbed by a chemical substance in order for a photochemical reaction to take place.
According to the second law of photochemistry, known as the Stark-Einstein law, for each photon of light absorbed by a chemical system, no more than one molecule is activated for a photochemical reaction, as defined by the quantum yield. When a molecule or atom in the ground state absorbs light, one electron is excited to a higher orbital level; this electron maintains its spin according to the spin selection rule. The excitation to a higher singlet state can be from HOMO to LUMO or to a higher orbital, so that singlet excitation states S1, S2, S3… at different energies are possible. Kasha's rule stipulates that higher singlet states would relax by radiationless decay or internal conversion to S1. Thus, S1 is but not always, the only relevant singlet excited state; this excited state S1 can further relax to S0 by IC, but by an allowed radiative transition from S1 to S0 that emits a photon. Alternatively, it is possible for the excited state S1 to undergo spin inversion and to generate a triplet excited state T1 having two unpaired electrons with the same spin.
This violation of the spin selection rule is possible by intersystem crossing of the vibrational and electronic levels of S1 and T1. According to Hund's rule of maximum multiplicity, this T1 state would be somewhat more stable than S1; this triplet state can relax to the ground state S0 by radiationless IC or by a radiation pathway called phosphorescence. This process implies a change of electronic spin, forbidden by spin selection rules, making phosphorescence much slower than fluorescence. Thus, triplet states have longer lifetimes than singlet states; these transitions are summarized in a state energy diagram or Jablonski diagram, the paradigm of molecular photochemistry. These excited species, either S1 or T1, have a half empty low-energy orbital, are more oxidizing than the ground state, but at the same time, they have an electron in a high energy orbital, are thus more reducing. In general, excited species are prone to participate in electron transfer processes. Photochemical reactions require a light source that emits wavelengths corresponding to an electronic transition in the reactant.
In the early experiments, sunlight was the light source. Mercury-vapor lamps are more common in the laboratory. Low pressure mercury vapor lamps emit at 254 nm. For polychromatic sources, wavelength ranges can be selected using filters. Alternatively, laser beams are monochromatic and LEDs have a narrowband that can be efficiently used, as well as Rayonet lamps, to get monochromatic beams; the emitted light must of course reach the targeted functional group without being blocked by the reactor, medium, or other functional groups present. For many applications, quartz is used for the reactors as well as to contain the lamp. Pyrex absorbs at wavelengths shorter than 275 nm; the solvent is an important experimental parameter. Solvents are potential reactants and for this reason, chlorinated solvents are avoided because the C-Cl bond can lead to chlorination of the substrate. Absorbing solvents prevent photons from reaching the substrate. Hydrocarbon solvents absorb only at short wavelengths and are thus preferred for photochemical experiments requiring high energy photons.
Solvents containing unsaturation absorb at longer wavelengths and can usefully filter out short wavelengths. For example and acetone "cut off" at wavelengths shorter than 215 and 330 nm, respectively. Continuous flow photochemistry offers multiple advantages over batch photochemistry. Photochemical reactions are driven by the number of photons that are able to activate molecules causing the desired reaction; the large surface area to volume ratio of a microreactor maximizes the illumination, at the same time allows for efficient cooling, which decreases the thermal side products. In the case of photochemical reactions, light provides the activation energy. Simplistically, light is one mechanism for providing the activation energy required for many reactions. If laser light is employed, it is possible to selectively excite a molecule so as to produce a desired electronic and vibrational state; the emission from a particular state may be selectively monitored, providing a measure of the population of that state
Lonsdaleite called hexagonal diamond in reference to the crystal structure, is an allotrope of carbon with a hexagonal lattice. In nature, it forms; the great heat and stress of the impact transforms the graphite into diamond, but retains graphite's hexagonal crystal lattice. Lonsdaleite was first identified in 1967 from the Canyon Diablo meteorite, where it occurs as microscopic crystals associated with diamond. Hexagonal diamond has been synthesized in the laboratory by compressing and heating graphite either in a static press or using explosives, it has been produced by chemical vapor deposition, by the thermal decomposition of a polymer, poly, at atmospheric pressure, under argon atmosphere, at 1,000 °C. It is translucent, brownish-yellow, has an index of refraction of 2.40 to 2.41 and a specific gravity of 3.2 to 3.3. Its hardness is theoretically superior to that of cubic diamond, according to computational simulations, but natural specimens exhibited somewhat lower hardness through a large range of values.
The cause is speculated as being due to the samples having been riddled with lattice defects and impurities. The property of lonsdaleite as a discrete material has been questioned, since specimens under crystallographic inspection showed not a bulk hexagonal lattice, but instead cubic diamond dominated by structural defects that include hexagonal sequences. A quantitative analysis of the X-ray diffraction data of lonsdaleite has shown that about equal amounts of hexagonal and cubic stacking sequences are present, it has been suggested that "stacking disordered diamond" is the most accurate structural description of lonsdaleite. On the other hand, recent shock experiments with in situ X-ray diffraction show strong evidence for creation of pure lonsdaleite in dynamic high-pressure environments such as meteorite impacts. According to the traditional picture, Lonsdaleite has a hexagonal unit cell, related to the diamond unit cell in the same way that the hexagonal and cubic close packed crystal systems are related.
The diamond structure can be considered to be made up of interlocking rings of six carbon atoms, in the chair conformation. In lonsdaleite, some rings are in the boat conformation instead. At the nanoscale dimensions cubic diamond is represented by diamondoids while hexagonal diamond is represented by wurtzoids. In diamond, all the carbon-to-carbon bonds, both within a layer of rings and between them, are in the staggered conformation, thus causing all four cubic-diagonal directions to be equivalent. Lonsdaleite is simulated to be 58% harder than diamond on the <100> face and to resist indentation pressures of 152 GPa, whereas diamond would break at 97 GPa. This is yet exceeded by IIa diamond's <111> tip hardness of 162 GPa. Lonsdaleite occurs as microscopic crystals associated with diamond in several meteorites: Canyon Diablo and Allan Hills 77283, it is naturally occurring in non-bolide diamond placer deposits in the Sakha Republic. Material with d-spacings consistent with Lonsdaleite has been found in sediments with uncertain dates at Lake Cuitzeo, in the state of Guanajuato, Mexico, by proponents of the controversial Younger Dryas impact hypothesis.
Its presence in local peat deposits is claimed as evidence for the Tunguska event being caused by a meteor rather than by a cometary fragment. Aggregated diamond nanorod Glossary of meteoritics List of minerals List of minerals named after people Anthony, J. W.. Mineralogy of Arizona. Tucson: University of Arizona Press. ISBN 0-8165-1579-4.. Mindat.org accessed 13 March 2005. Webmineral accessed 13 March 2005. Materials Science and Technology Division, Naval Research Laboratory website accessed 14 May 2006. Diamond no longer nature's hardest material lonsdaleite 3D animation
Jöns Jacob Berzelius
Baron Jöns Jacob Berzelius, named by himself and contemporary society as Jacob Berzelius, was a Swedish chemist. Berzelius is considered, along with Robert Boyle, John Dalton, Antoine Lavoisier, to be one of the founders of modern chemistry. Berzelius began his career as a physician but his researches in physical chemistry were of lasting significance in the development of the subject, he is noted for his determination of atomic weights. In 1803 Berzelius demonstrated the power of an electrochemical cell to decompose chemicals into pairs of electrically opposite constituents. Berzelius's work with atomic weights and his theory of electrochemical dualism led to his development of a modern system of chemical formula notation that could portray the composition of any compound both qualitatively and quantitatively, his system abbreviated the Latin names of the elements with one or two letters and applied subscripts to designate the number of atoms of each element present in both the acidic and basic ingredients.
Berzelius himself isolated several new elements, including cerium and thorium. Berzelius’s interest in mineralogy fostered his analysis and preparation of new compounds of these and other elements; the mineral berzelianite was named after him. He was a strict empiricist and insisted that any new theory be consistent with the sum of chemical knowledge, he developed classical analytical techniques, investigated isomerism and catalysis, phenomena that owe their names to him. He became a member of the Royal Swedish Academy of Sciences in 1808 and served from 1818 as its principal functionary, the perpetual secretary, he is known in Sweden as "the Father of Swedish Chemistry". Berzelius Day is celebrated on 20 August in honour of him. Berzelius was born in the parish of Väversunda in Östergötland in Sweden, his father was a school teacher in his mother a homemaker. Berzelius lost both his parents at an early age. Relatives in Linköping took care of him, there he attended the school today known as Katedralskolan.
He enrolled at Uppsala University, where he learned the profession of medical doctor from 1796 to 1801. He worked as an apprentice with a physician in the Medevi mineral springs. During this time, he conducted analysis of the spring water. For his medical studies, he examined the influence of galvanic current on several diseases and graduated as M. D. in 1802. He worked as physician near Stockholm until the mine-owner Wilhelm Hisinger discovered his analytical abilities and provided him with a laboratory. Between 1808 and 1836, Berzelius worked together with Anna Sundström. In 1807, Berzelius was appointed professor in pharmacy at the Karolinska Institute. In 1808, he was elected a member of the Royal Swedish Academy of Sciences. At this time, the Academy had been stagnating for several years, since the era of romanticism in Sweden had led to less interest in the sciences. In 1818, Berzelius was elected the Academy's secretary and held the post until 1848. During Berzelius' tenure, he is credited with revitalising the Academy and bringing it into a second golden era.
He was elected a Foreign Honorary Member of the American Academy of Arts and Sciences in 1822. In 1827, he became correspondent of the Royal Institute of the Netherlands, in 1830 associate member. In 1837, he was elected a member of the Swedish Academy, on chair number 5. Not long after arriving to Stockholm he wrote a chemistry textbook for his medical students, from which point a long and fruitful career in chemistry began. In 1813, he published an essay on the proportions of elements in compounds; the essay commenced with a general description, introduced his new symbolism, examined all the known elements, included a table of specific weights, finished with a selection of compounds written in his new formalisation. In 1818, he compiled a table of relative atomic weights, where oxygen was set to 100, which included all of the elements known at the time; this work provided evidence in favour of the atomic theory proposed by John Dalton: that inorganic chemical compounds are composed of atoms combined in whole number amounts.
In discovering that atomic weights are not integer multiples of the weight of hydrogen, Berzelius disproved Prout's hypothesis that elements are built up from atoms of hydrogen. Berzelius's atomic weight tables was first published in a German translation of his Textbook of Chemistry in 1826. In order to aid his experiments, he developed a system of chemical notation in which the elements were given simple written labels—such as O for oxygen, or Fe for iron—with proportions noted by numbers; this is the same system used today, the only difference being that instead of the subscript number used today, Berzelius used a superscript. Berzelius is credited with identifying the chemical elements silicon, selenium and cerium. Students working in Berzelius's laboratory discovered lithium and vanadium. Berzelius discovered silicon by repeating an experiment performed by Thénard. In the experiment, Berzelius reacted silicon tetrafluoride with potassium metal and purified its product by washing it until it became a brown powde
Graphite, archaically referred to as plumbago, is a crystalline form of the element carbon with its atoms arranged in a hexagonal structure. It occurs in this form and is the most stable form of carbon under standard conditions. Under high pressures and temperatures it converts to diamond. Graphite is used in lubricants, its high conductivity makes it useful in electronic products such as electrodes and solar panels. The principal types of natural graphite, each occurring in different types of ore deposits, are Crystalline small flakes of graphite occurs as isolated, plate-like particles with hexagonal edges if unbroken; when broken the edges can be angular. Ordered pyrolytic graphite refers to graphite with an angular spread between the graphite sheets of less than 1°; the name "graphite fiber" is sometimes used to refer to carbon fibers or carbon fiber-reinforced polymer. Graphite occurs in metamorphic rocks as a result of the reduction of sedimentary carbon compounds during metamorphism, it occurs in igneous rocks and in meteorites.
Minerals associated with graphite include quartz, calcite and tourmaline. The principal export sources of mined graphite are in order of tonnage: China, Canada and Madagascar. In meteorites, graphite occurs with silicate minerals. Small graphitic crystals in meteoritic iron are called cliftonite; some microscopic grains have distinctive isotopic compositions, indicating that they were formed before the Solar system. They are one of about 12 known types of mineral that predate the Solar System and have been detected in molecular clouds; these minerals were formed in the ejecta when supernovae exploded or low- to intermediate-sized stars expelled their outer envelopes late in their lives. Graphite may be the third oldest mineral in the Universe. Solid carbon comes in different forms known as allotropes depending on the type of chemical bond; the two most common are graphite. In diamond the bonds are sp3 and the atoms form tetrahedra with each bound to four nearest neighbors. In graphite they are sp2 orbital hybrids and the atoms form in planes with each bound to three nearest neighbors 120 degrees apart.
The individual layers are called graphene. In each layer, the carbon atoms are arranged in a honeycomb lattice with separation of 0.142 nm, the distance between planes is 0.335 nm. Atoms in the plane are bonded covalently, with only three of the four potential bonding sites satisfied; the fourth electron is free to migrate in the plane. However, it does not conduct in a direction at right angles to the plane. Bonding between layers is via weak van der Waals bonds, which allows layers of graphite to be separated, or to slide past each other; the two known forms of graphite and beta, have similar physical properties, except that the graphene layers stack differently. The alpha graphite may be either buckled; the alpha form can be converted to the beta form through mechanical treatment and the beta form reverts to the alpha form when it is heated above 1300 °C. The equilibrium pressure and temperature conditions for a transition between graphite and diamond is well established theoretically and experimentally.
The pressure changes linearly between 1.7 GPa at 0 K and 12 GPa at 5000 K. However, the phases have a wide region about this line where they can coexist. At normal temperature and pressure, 20 °C and 1 standard atmosphere, the stable phase of carbon is graphite, but diamond is metastable and its rate of conversion to graphite is negligible. However, at temperatures above about 4500 K, diamond converts to graphite. Rapid conversion of graphite to diamond requires pressures well above the equilibrium line: at 2000 K, a pressure of 35 GPa is needed; the acoustic and thermal properties of graphite are anisotropic, since phonons propagate along the bound planes, but are slower to travel from one plane to another. Graphite's high thermal stability and electrical and thermal conductivity facilitate its widespread use as electrodes and refractories in high temperature material processing applications. However, in oxygen-containing atmospheres graphite oxidizes to form carbon dioxide at temperatures of 700 °C and above.
Graphite is hence useful in such applications as arc lamp electrodes. It can conduct electricity due to the vast electron delocalization within the carbon layers; these valence electrons are free to move. However, the electricity is conducted within the plane of the layers; the conductive properties of powdered graphite allow its use as pressure sensor in carbon microphones. Graphite and graphite powder are valued in industrial applications for their self-lubricating and dry lubricating properties. There is a common belief that graphite's lubricating properties are due to the loose interlamellar coupling between sheets in the structure. However, it has been shown that in a vacuum environment, graphite degrades as a lubricant, due to the hypoxic conditions; this observation led to the hypothesis that the lubrication is due to the presence of fluids between the layers, such as air and water, which are adsorbed from the