In quantum mechanics and particle physics, spin is an intrinsic form of angular momentum carried by elementary particles, composite particles, atomic nuclei. Spin is one of two types of angular momentum in quantum mechanics, the other being orbital angular momentum; the orbital angular momentum operator is the quantum-mechanical counterpart to the classical angular momentum of orbital revolution and appears when there is periodic structure to its wavefunction as the angle varies. The existence of spin angular momentum is inferred from experiments, such as the Stern–Gerlach experiment, in which silver atoms were observed to possess two possible discrete angular momenta despite having no orbital angular momentum. In some ways, spin is like a vector quantity. All elementary particles of a given kind have the same magnitude of spin angular momentum, indicated by assigning the particle a spin quantum number; the SI unit of spin is the or, just as with classical angular momentum. In practice, spin is given as a dimensionless spin quantum number by dividing the spin angular momentum by the reduced Planck constant ħ, which has the same units of angular momentum, although this is not the full computation of this value.
The "spin quantum number" is called "spin", leaving its meaning as the unitless "spin quantum number" to be inferred from context. When combined with the spin-statistics theorem, the spin of electrons results in the Pauli exclusion principle, which in turn underlies the periodic table of chemical elements. Wolfgang Pauli in 1924 was the first to propose a doubling of electron states due to a two-valued non-classical "hidden rotation". In 1925, George Uhlenbeck and Samuel Goudsmit at Leiden University suggested the simple physical interpretation of a particle spinning around its own axis, in the spirit of the old quantum theory of Bohr and Sommerfeld. Ralph Kronig anticipated the Uhlenbeck-Goudsmit model in discussion with Hendrik Kramers several months earlier in Copenhagen, but did not publish; the mathematical theory was worked out in depth by Pauli in 1927. When Paul Dirac derived his relativistic quantum mechanics in 1928, electron spin was an essential part of it; as the name suggests, spin was conceived as the rotation of a particle around some axis.
This picture is correct so far as spin obeys the same mathematical laws as quantized angular momenta do. On the other hand, spin has some peculiar properties that distinguish it from orbital angular momenta: Spin quantum numbers may take half-integer values. Although the direction of its spin can be changed, an elementary particle cannot be made to spin faster or slower; the spin of a charged particle is associated with a magnetic dipole moment with a g-factor differing from 1. This could only occur classically if the internal charge of the particle were distributed differently from its mass; the conventional definition of the spin quantum number, s, is s = n/2, where n can be any non-negative integer. Hence the allowed values of s are 1/2, 1, 3/2, 2, etc.. The value of s for an elementary particle depends only on the type of particle, cannot be altered in any known way; the spin angular momentum, S, of any physical system is quantized. The allowed values of S are S = ℏ s = h 4 π n, where h is the Planck constant and ℏ = h/2π is the reduced Planck constant.
In contrast, orbital angular momentum can only take on integer values of s. Those particles with half-integer spins, such as 1/2, 3/2, 5/2, are known as fermions, while those particles with integer spins, such as 0, 1, 2, are known as bosons; the two families of particles obey different rules and broadly have different roles in the world around us. A key distinction between the two families is. In contrast, bosons obey the rules of Bose–Einstein statistics and have no such restriction, so they may "bunch together" if in identical states. Composite particles can have spins different from their component particles. For example, a helium atom in the ground state has spin 0 and behaves like a boson though the quarks and electrons which make it up are all fermions; this has some profound consequences: Quarks and leptons, which make up what is classically known as matter, are all fermions with spin 1/2. The common idea that "matter takes up space" comes from the Pauli exclusion principle acting on these particles to prevent the fermions that make up matter from being in the same quantum state.
Further compaction would require electrons to occupy the same energy states, therefore a kind of pressure acts to resist the fermions being overly close. Elementary fermions with other spins are not known to exist. Elementary particles which are thought of as carrying forces are all bosons with spin 1, they include the photon which carries the electromagnetic force, the gluon, the W and Z bosons. The ability of bosons to occupy the same quantu
Proteomics is the large-scale study of proteins. Proteins are vital parts with many functions; the term proteomics was coined in analogy to genomics, the study of the genome. The word proteome is a portmanteau of protein and genome, was coined by Marc Wilkins in 1994 while he was a Ph. D. student at Macquarie University. Macquarie University founded the first dedicated proteomics laboratory in 1995; the proteome is the entire set of proteins, produced or modified by an organism or system. Proteomics has enabled the identification of increasing numbers of protein; this stresses, that a cell or organism undergoes. Proteomics is an interdisciplinary domain that has benefitted from the genetic information of various genome projects, including the Human Genome Project, it covers the exploration of proteomes from the overall level of protein composition and activity. It is an important component of functional genomics. Proteomics refers to the large-scale experimental analysis of proteins and proteomes, but is used to refer to protein purification and mass spectrometry.
After genomics and transcriptomics, proteomics is the next step in the study of biological systems. It is more complicated than genomics because an organism's genome is more or less constant, whereas proteomes differ from cell to cell and from time to time. Distinct genes are expressed in different cell types, which means that the basic set of proteins that are produced in a cell needs to be identified. In the past this phenomenon was assessed by RNA analysis, but it was found to lack correlation with protein content. Now it is known that mRNA is not always translated into protein, the amount of protein produced for a given amount of mRNA depends on the gene it is transcribed from and on the current physiological state of the cell. Proteomics confirms the presence of the protein and provides a direct measure of the quantity present. Not only does the translation from mRNA cause differences, but many proteins are subjected to a wide variety of chemical modifications after translation; the most common and studied post translational modifications include phosphorylation and glycosylation.
Many of these post-translational modifications are critical to the protein's function. One such modification is phosphorylation, which happens to many enzymes and structural proteins in the process of cell signaling; the addition of a phosphate to particular amino acids—most serine and threonine mediated by serine-threonine kinases, or more tyrosine mediated by tyrosine kinases—causes a protein to become a target for binding or interacting with a distinct set of other proteins that recognize the phosphorylated domain. Because protein phosphorylation is one of the most-studied protein modifications, many "proteomic" efforts are geared to determining the set of phosphorylated proteins in a particular cell or tissue-type under particular circumstances; this alerts the scientist to the signaling pathways. Ubiquitin is a small protein that may be affixed to certain protein substrates by enzymes called E3 ubiquitin ligases. Determining which proteins are poly-ubiquitinated helps understand how protein pathways are regulated.
This is, therefore, an additional legitimate "proteomic" study. Once a researcher determines which substrates are ubiquitinated by each ligase, determining the set of ligases expressed in a particular cell type is helpful. In addition to phosphorylation and ubiquitination, proteins may be subjected to methylation, glycosylation and nitrosylation; some proteins undergo all these modifications in time-dependent combinations. This illustrates the potential complexity of studying protein function. A cell may make different sets of proteins at different times or under different conditions, for example during development, cellular differentiation, cell cycle, or carcinogenesis. Further increasing proteome complexity, as mentioned, most proteins are able to undergo a wide range of post-translational modifications. Therefore, a "proteomics" study may become complex quickly if the topic of study is restricted. In more ambitious settings, such as when a biomarker for a specific cancer subtype is sought, the proteomics scientist might elect to study multiple blood serum samples from multiple cancer patients to minimise confounding factors and account for experimental noise.
Thus, complicated experimental designs are sometimes necessary to account for the dynamic complexity of the proteome. Proteomics gives a different level of understanding than genomics for many reasons: the level of transcription of a gene gives only a rough estimate of its level of translation into a protein. An mRNA produced in abundance may be degraded or translated inefficiently, resulting in a small amount of protein; as mentioned above many proteins experience post-translational modifications that profoundly affect their activities. Methods such as phosphoproteomics and glycoproteomics are used to study post-translational modifications. Many transcripts give rise to more than one protein, through alternative splicing or alternative post-translational modifications. Many proteins form complexes with other proteins or RNA molecules, only function in the presence of these other molecules. Protein degradation rate plays an important role in protein content. Reproducibility. One major factor affecting reproducibility in proteomics experiments is the simultaneous elution of many more peptides than mass spectrometers can measure.
This causes stochastic differences between experiments due to dat
Isotope-ratio mass spectrometry
Isotope-ratio mass spectrometry is a specialization of mass spectrometry, in which mass spectrometric methods are used to measure the relative abundance of isotopes in a given sample. This technique has two different applications in environmental sciences; the analysis of'stable isotopes' is concerned with measuring isotopic variations arising from mass-dependent isotopic fractionation in natural systems. On the other hand, radiogenic isotope analysis involves measuring the abundances of decay-products of natural radioactivity, is used in most long-lived radiometric dating methods; the isotope-ratio mass spectrometer allows the precise measurement of mixtures of occurring isotopes. Most instruments used for precise determination of isotope ratios are of the magnetic sector type; this type of analyzer is superior to the quadrupole type in this field of research for two reasons. First, it can be set up for multiple-collector analysis, second, it gives high-quality'peak shapes'. Both of these considerations are important for isotope-ratio analysis at high precision and accuracy.
The sector-type instrument designed by Alfred Nier was such an advance in mass spectrometer design that this type of instrument is called the'Nier type'. In the most general terms the instrument operates by ionizing the sample of interest, accelerating it over a potential in the kilo-volt range, separating the resulting stream of ions according to their mass-to-charge ratio. Beams with lighter ions bend at a smaller radius; the current of each ion beam is measured using a'Faraday cup' or multiplier detector. Many radiogenic isotope measurements are made by ionization of a solid source, whereas stable isotope measurements of light elements are made in an instrument with a gas source. In a "multicollector" instrument, the ion collector has an array of Faraday cups, which allows the simultaneous detection of multiple isotopes. Measurement of natural variations in the abundances of stable isotopes of the same element is referred to as stable isotope analysis; this field is of interest because the differences in mass between different isotopes leads to isotope fractionation, causing measurable effects on the isotopic composition of samples, characteristic of their biological or physical history.
As a specific example, the hydrogen isotope deuterium is double the mass of the common hydrogen isotope. Water molecules containing the common hydrogen isotope have a mass of 18. Water incorporating a deuterium atom has a mass of 19, over 5% heavier; the energy to vaporise the heavy water molecule is higher than that to vaporize the normal water so isotope fractionation occurs during the process of evaporation. Thus a sample of sea water will exhibit a quite detectable isotopic-ratio difference when compared to Antarctic snowfall. Samples must be introduced to the mass spectrometer as pure gases, achieved through combustion, gas chromatographic feeds, or chemical trapping. By comparing the detected isotopic ratios to a measured standard, an accurate determination of the isotopic make up of the sample is obtained. For example, carbon isotope ratios are measured relative to the international standard for C; the C standard is produced from a fossil belemnite found in the Peedee Formation, a limestone formed in the Cretaceous period in South Carolina, U.
S. A; the fossil is referred to as VPDB and has 13C:12C ratio of 0.0112372. Oxygen isotope ratios are measured relative the standard, V-SMOW, it is critical that the sample be processed before entering the mass spectrometer so that only a single chemical species enters at a given time. Samples are combusted or pyrolyzed and the desired gas species is purified by means of traps, catalysts and/or chromatography; the two most common types of IRMS instruments are dual inlet. In dual inlet IRMS, purified gas obtained from a sample is alternated with a standard gas by means of a system of valves, so that a number of comparison measurements are made of both gases. In continuous flow IRMS, sample preparation occurs before introduction to the IRMS, the purified gas produced from the sample is measured just once; the standard gas may be measured before and after the sample or after a series of sample measurements. While continuous-flow IRMS instruments can achieve higher sample throughput and are more convenient to use than dual inlet instruments, the yielded data is of 10-fold lower precision.
A static gas mass spectrometer is one in which a gaseous sample for analysis is fed into the source of the instrument and left in the source without further supply or pumping throughout the analysis. This method can be used for'stable isotope' analysis of light gases, but it is used in the isotopic analysis of noble gases for radiometric dating or isotope geochemistry. Important examples are argon -- argon helium isotope analysis. Several of the isotope systems involved in radiometric dating depend on IRMS using thermal ionization of a solid sample loaded into the source of the mass spectrometer; these methods include rubidium–strontium dating, uranium–lead dating, lead–lead dating and samarium–neodymium dating. When these isotope ratios are measured by TIMS, mass-dependent fractionation occurs as species are emitted by the hot filament. Fractionation occurs due to the excitation of the sample and therefore must
Isotopes of boron
Boron occurs as isotopes 10B and 11B, the latter of which makes up about 80% of natural boron. There are 13 radioisotopes that have been discovered, with mass numbers from 7 to 21, all with short half-lives, the longest being that of 8B, with a half-life of only 770 milliseconds and 12B with a half-life of 20.2 ms. All other isotopes have half-lives shorter than 17.35 ms. Those isotopes with mass below 10 decay into helium while those with mass above 11 become carbon; the precision of the isotope abundances and atomic mass is limited through variations. The given ranges should be applicable to any normal terrestrial material. Commercially available materials may have been subjected to an undisclosed or inadvertent isotopic fractionation. Substantial deviations from the given mass and composition can occur. Values marked # are not purely derived from experimental data, but at least from systematic trends. Spins with weak assignment arguments are enclosed in parentheses. Uncertainties are given in concise form in parentheses after the corresponding last digits.
Uncertainty values denote one standard deviation, except isotopic composition and standard atomic mass from IUPAC, which use folical uncertainties. Nuclide masses are given by IUPAP Commission on Symbols, Nomenclature, Atomic Masses and Fundamental Constants. Isotope abundances are given by IUPAC Commission on Atomic Weights. Neutrinos from boron-8 beta decays within the sun are an important background to dark matter direct detection experiments, they are the first component of the neutrino floor that dark matter direct detection experiments are expected to encounter. Boron-10 is used in boron neutron capture therapy as an experimental treatment of some brain cancers. Isotope masses from: Audi, Georges. P.. "Atomic weights of the elements. Review 2000". Pure and Applied Chemistry. 75: 683–800. Doi:10.1351/pac200375060683. Wieser, Michael E.. "Atomic weights of the elements 2005". Pure and Applied Chemistry. 78: 2051–2066. Doi:10.1351/pac200678112051. Lay summary. Half-life and isomer data selected from the following sources.
See editing notes on this article's talk page. Audi, Georges. "NuDat 2.x database". Brookhaven National Laboratory. Holden, Norman E.. "11. Table of the Isotopes". In Lide, David R. CRC Handbook of Physics. Boca Raton, Florida: CRC Press. ISBN 978-0-8493-0485-9
The atomic number or proton number of a chemical element is the number of protons found in the nucleus of an atom. It is identical to the charge number of the nucleus; the atomic number uniquely identifies a chemical element. In an uncharged atom, the atomic number is equal to the number of electrons; the sum of the atomic number Z and the number of neutrons, N, gives the mass number A of an atom. Since protons and neutrons have the same mass and the mass defect of nucleon binding is always small compared to the nucleon mass, the atomic mass of any atom, when expressed in unified atomic mass units, is within 1% of the whole number A. Atoms with the same atomic number Z but different neutron numbers N, hence different atomic masses, are known as isotopes. A little more than three-quarters of occurring elements exist as a mixture of isotopes, the average isotopic mass of an isotopic mixture for an element in a defined environment on Earth, determines the element's standard atomic weight, it was these atomic weights of elements that were the quantities measurable by chemists in the 19th century.
The conventional symbol Z comes from the German word Zahl meaning number, before the modern synthesis of ideas from chemistry and physics denoted an element's numerical place in the periodic table, whose order is but not consistent with the order of the elements by atomic weights. Only after 1915, with the suggestion and evidence that this Z number was the nuclear charge and a physical characteristic of atoms, did the word Atomzahl come into common use in this context. Loosely speaking, the existence or construction of a periodic table of elements creates an ordering of the elements, so they can be numbered in order. Dmitri Mendeleev claimed. However, in consideration of the elements' observed chemical properties, he changed the order and placed tellurium ahead of iodine; this placement is consistent with the modern practice of ordering the elements by proton number, Z, but that number was not known or suspected at the time. A simple numbering based on periodic table position was never satisfactory, however.
Besides the case of iodine and tellurium several other pairs of elements were known to have nearly identical or reversed atomic weights, thus requiring their placement in the periodic table to be determined by their chemical properties. However the gradual identification of more and more chemically similar lanthanide elements, whose atomic number was not obvious, led to inconsistency and uncertainty in the periodic numbering of elements at least from lutetium onward. In 1911, Ernest Rutherford gave a model of the atom in which a central core held most of the atom's mass and a positive charge which, in units of the electron's charge, was to be equal to half of the atom's atomic weight, expressed in numbers of hydrogen atoms; this central charge would thus be half the atomic weight. In spite of Rutherford's estimation that gold had a central charge of about 100, a month after Rutherford's paper appeared, Antonius van den Broek first formally suggested that the central charge and number of electrons in an atom was equal to its place in the periodic table.
This proved to be the case. The experimental position improved after research by Henry Moseley in 1913. Moseley, after discussions with Bohr, at the same lab, decided to test Van den Broek's and Bohr's hypothesis directly, by seeing if spectral lines emitted from excited atoms fitted the Bohr theory's postulation that the frequency of the spectral lines be proportional to the square of Z. To do this, Moseley measured the wavelengths of the innermost photon transitions produced by the elements from aluminum to gold used as a series of movable anodic targets inside an x-ray tube; the square root of the frequency of these photons increased from one target to the next in an arithmetic progression. This led to the conclusion that the atomic number does correspond to the calculated electric charge of the nucleus, i.e. the element number Z. Among other things, Moseley demonstrated that the lanthanide series must have 15 members—no fewer and no more—which was far from obvious from the chemistry at that time.
After Moseley's death in 1915, the atomic numbers of all known elements from hydrogen to uranium were examined by his method. There were seven elements which were not found and therefore identified as still undiscovered, corresponding to atomic numbers 43, 61, 72, 75, 85, 87 and 91. From 1918 to 1947, all seven of these missing elements were discovered. By this time the first four transuranium elements had been discovered, so that the periodic table was complete with no gaps as far as curium. In 1915 the rea
The environmental isotopes are a subset of the isotopes, both stable and radioactive, which are the object of isotope geochemistry. They are used as tracers to see how things move around within the ocean-atmosphere system, within terrestrial biomes, within the Earth's surface, between these broad domains. Chemical elements are defined by their number of protons, but the mass of the atom is determined by the number of protons and neutrons in the nucleus. Isotopes are atoms that are of a specific element, but have different numbers of neutrons and thus different masses. In a specific object, you can have a ratio between two isotopes of an element; this ratio varies in the world, so in order to study isotopic ratio changes across the world, changes in isotope ratios are defined as deviations from a standard, multiplied by 1000. This unit is a "per mil"; as a convention, the ratio is of the heavier isotope to the lower isotope. Δ C 13 = × 1000 ‰. They are classified as mass independent fractionation and mass dependent fractionation.
An example of a mass independent process is the fractionation of oxygen atoms in ozone. This is due to the kinetic isotope effect and is caused by different isotope molecules reacting at different speeds. An example of a mass dependent process is the fractionation of water as it transitions from the liquid to gas phase. Water molecules with heavier isotopes tend to stay in the liquid phase as water molecules with lighter isotopes preferentially move to the gas phase. Of the different isotopes that exist, one common classification is distinguishing radioactive isotopes from stable isotopes. Radioactive isotopes are isotopes. For example, 3H is a radioactive isotope of hydrogen, it decays into 3He with a half-life of ~12.3 years. By comparison, stable isotopes are more stable, decaying much more and having much longer half-lives. Examples of stable isotopes are 87Sr; these isotopes of strontium have half-lives on the order of billions of years or are unmeasured because of how stable they are. On timescales that geologists and environmental scientists investigate, these isotopes are stable.
Both of these types of isotopes are useful to scientists. Radioactive isotopes are more useful on shorter timescales, such as investigating modern circulation of the ocean using 14C, while stable isotopes are more useful on longer timescales, such as investigating differences in river flow with strontium stable isotopes; these isotopes are used as tracers to study various phenomena of interest. These tracers have a certain distribution spatially, so scientists need to deconvolve the different processes that affect these tracer distributions. One way tracer distributions are set is by conservative mixing. In conservative mixing, the amount of the tracer is conserved. An example of this is mixing two water masses with different salinities; the salt from the saltier water mass moves to the less salty water mass, keeping the total amount of salinity constant. This way of mixing tracers is important, giving a baseline of what value of a tracer one should expect; the value of a tracer as a point is expected to be an average value of the sources that flow into that region.
Deviations from this are indicative of other processes. These can be called nonconservative mixing, where there are other processes that do not conserve the amount of tracer. An example of this is 14C; this mixes between water masses, but it decays over time, reducing the amount of 14C in the region. The most used environmental isotopes are: deuterium tritium carbon-13 carbon-14 nitrogen-15 oxygen-18 silicon-29 chlorine-36 isotopes of uranium isotopes of strontium One topic that environmental isotopes are used to study is the circulation of the ocean. Treating the ocean as a box is only useful in some studies; this leads to an understanding of. This helps deconvolve circulation effects from other phenomena that affect certain tracers such as radioactive and biological processes. Using rudimentary observation techniques, the circulation of the surface ocean can be determined. In the Atlantic basin, surface waters flow from the south towards the north in general, while creating gyres in the northern and southern Atlantic.
In the Pacific Ocean, the gyres still form, but there is comparatively little large scale meridional movement. For deep waters, there are two areas; these are in the Antarctic. The deep water masses formed are North Atla
Carbon is a chemical element with symbol C and atomic number 6. It is nonmetallic and tetravalent—making four electrons available to form covalent chemical bonds, it belongs to group 14 of the periodic table. Three isotopes occur 12C and 13C being stable, while 14C is a radionuclide, decaying with a half-life of about 5,730 years. Carbon is one of the few elements known since antiquity. Carbon is the 15th most abundant element in the Earth's crust, the fourth most abundant element in the universe by mass after hydrogen and oxygen. Carbon's abundance, its unique diversity of organic compounds, its unusual ability to form polymers at the temperatures encountered on Earth enables this element to serve as a common element of all known life, it is the second most abundant element in the human body by mass after oxygen. The atoms of carbon can bond together in different ways, termed allotropes of carbon; the best known are graphite and amorphous carbon. The physical properties of carbon vary with the allotropic form.
For example, graphite is opaque and black while diamond is transparent. Graphite is soft enough to form a streak on paper, while diamond is the hardest occurring material known. Graphite is a good electrical conductor. Under normal conditions, carbon nanotubes, graphene have the highest thermal conductivities of all known materials. All carbon allotropes are solids under normal conditions, with graphite being the most thermodynamically stable form at standard temperature and pressure, they are chemically resistant and require high temperature to react with oxygen. The most common oxidation state of carbon in inorganic compounds is +4, while +2 is found in carbon monoxide and transition metal carbonyl complexes; the largest sources of inorganic carbon are limestones and carbon dioxide, but significant quantities occur in organic deposits of coal, peat and methane clathrates. Carbon forms a vast number of compounds, more than any other element, with ten million compounds described to date, yet that number is but a fraction of the number of theoretically possible compounds under standard conditions.
For this reason, carbon has been referred to as the "king of the elements". The allotropes of carbon include graphite, one of the softest known substances, diamond, the hardest occurring substance, it bonds with other small atoms, including other carbon atoms, is capable of forming multiple stable covalent bonds with suitable multivalent atoms. Carbon is known to form ten million different compounds, a large majority of all chemical compounds. Carbon has the highest sublimation point of all elements. At atmospheric pressure it has no melting point, as its triple point is at 10.8±0.2 MPa and 4,600 ± 300 K, so it sublimes at about 3,900 K. Graphite is much more reactive than diamond at standard conditions, despite being more thermodynamically stable, as its delocalised pi system is much more vulnerable to attack. For example, graphite can be oxidised by hot concentrated nitric acid at standard conditions to mellitic acid, C66, which preserves the hexagonal units of graphite while breaking up the larger structure.
Carbon sublimes in a carbon arc, which has a temperature of about 5800 K. Thus, irrespective of its allotropic form, carbon remains solid at higher temperatures than the highest-melting-point metals such as tungsten or rhenium. Although thermodynamically prone to oxidation, carbon resists oxidation more than elements such as iron and copper, which are weaker reducing agents at room temperature. Carbon is the sixth element, with a ground-state electron configuration of 1s22s22p2, of which the four outer electrons are valence electrons, its first four ionisation energies, 1086.5, 2352.6, 4620.5 and 6222.7 kJ/mol, are much higher than those of the heavier group-14 elements. The electronegativity of carbon is 2.5 higher than the heavier group-14 elements, but close to most of the nearby nonmetals, as well as some of the second- and third-row transition metals. Carbon's covalent radii are taken as 77.2 pm, 66.7 pm and 60.3 pm, although these may vary depending on coordination number and what the carbon is bonded to.
In general, covalent radius decreases with higher bond order. Carbon compounds form the basis of all known life on Earth, the carbon–nitrogen cycle provides some of the energy produced by the Sun and other stars. Although it forms an extraordinary variety of compounds, most forms of carbon are comparatively unreactive under normal conditions. At standard temperature and pressure, it resists all but the strongest oxidizers, it does not react with hydrochloric acid, chlorine or any alkalis. At elevated temperatures, carbon reacts with oxygen to form carbon oxides and will rob oxygen from metal oxides to leave the elemental metal; this exothermic reaction is used in the iron and steel industry to smelt iron and to control the carbon content of steel: Fe3O4 + 4 C → 3 Fe + 4 COCarbon monoxide can be recycled to smelt more iron: Fe3O4 + 4 CO → 3 Fe + 4 CO2with sulfur to form carbon disulfide and with steam in the coal-gas reaction: C + H2O → CO + H2. Carbon combines with some metals at high temperatures to form metallic carbides, such as the iron carbide cementite in steel and tungsten carbide used as an abrasive and for making hard tips for cutting tools.
The system of carbon allotropes spans a range of extremes: Atomic carbon is a ver