A chemical element is a species of atom having the same number of protons in their atomic nuclei. For example, the atomic number of oxygen is 8, so the element oxygen consists of all atoms which have 8 protons. 118 elements have been identified, of which the first 94 occur on Earth with the remaining 24 being synthetic elements. There are 80 elements that have at least one stable isotope and 38 that have radionuclides, which decay over time into other elements. Iron is the most abundant element making up Earth, while oxygen is the most common element in the Earth's crust. Chemical elements constitute all of the ordinary matter of the universe; however astronomical observations suggest that ordinary observable matter makes up only about 15% of the matter in the universe: the remainder is dark matter. The two lightest elements and helium, were formed in the Big Bang and are the most common elements in the universe; the next three elements were formed by cosmic ray spallation, are thus rarer than heavier elements.
Formation of elements with from 6 to 26 protons occurred and continues to occur in main sequence stars via stellar nucleosynthesis. The high abundance of oxygen and iron on Earth reflects their common production in such stars. Elements with greater than 26 protons are formed by supernova nucleosynthesis in supernovae, when they explode, blast these elements as supernova remnants far into space, where they may become incorporated into planets when they are formed; the term "element" is used for atoms with a given number of protons as well as for a pure chemical substance consisting of a single element. For the second meaning, the terms "elementary substance" and "simple substance" have been suggested, but they have not gained much acceptance in English chemical literature, whereas in some other languages their equivalent is used. A single element can form multiple substances differing in their structure; when different elements are chemically combined, with the atoms held together by chemical bonds, they form chemical compounds.
Only a minority of elements are found uncombined as pure minerals. Among the more common of such native elements are copper, gold and sulfur. All but a few of the most inert elements, such as noble gases and noble metals, are found on Earth in chemically combined form, as chemical compounds. While about 32 of the chemical elements occur on Earth in native uncombined forms, most of these occur as mixtures. For example, atmospheric air is a mixture of nitrogen and argon, native solid elements occur in alloys, such as that of iron and nickel; the history of the discovery and use of the elements began with primitive human societies that found native elements like carbon, sulfur and gold. Civilizations extracted elemental copper, tin and iron from their ores by smelting, using charcoal. Alchemists and chemists subsequently identified many more; the properties of the chemical elements are summarized in the periodic table, which organizes the elements by increasing atomic number into rows in which the columns share recurring physical and chemical properties.
Save for unstable radioactive elements with short half-lives, all of the elements are available industrially, most of them in low degrees of impurities. The lightest chemical elements are hydrogen and helium, both created by Big Bang nucleosynthesis during the first 20 minutes of the universe in a ratio of around 3:1 by mass, along with tiny traces of the next two elements and beryllium. All other elements found in nature were made by various natural methods of nucleosynthesis. On Earth, small amounts of new atoms are produced in nucleogenic reactions, or in cosmogenic processes, such as cosmic ray spallation. New atoms are naturally produced on Earth as radiogenic daughter isotopes of ongoing radioactive decay processes such as alpha decay, beta decay, spontaneous fission, cluster decay, other rarer modes of decay. Of the 94 occurring elements, those with atomic numbers 1 through 82 each have at least one stable isotope. Isotopes considered stable are those. Elements with atomic numbers 83 through 94 are unstable to the point that radioactive decay of all isotopes can be detected.
Some of these elements, notably bismuth and uranium, have one or more isotopes with half-lives long enough to survive as remnants of the explosive stellar nucleosynthesis that produced the heavy metals before the formation of our Solar System. At over 1.9×1019 years, over a billion times longer than the current estimated age of the universe, bismuth-209 has the longest known alpha decay half-life of any occurring element, is always considered on par with the 80 stable elements. The heaviest elements undergo radioactive decay with half-lives so short that they are not found in nature and must be synthesized; as of 2010, there are 118 known elements (in this context, "known" means observed well enough from just a few de
The S3 molecule or trisulfur or sulfur trimer or thiozone or triatomic sulfur is an allotrope of sulfur. It occurs as a mixture in liquid and gaseous sulfur and at cryogenic temperatures as a solid. Under standard conditions it is unstable and self reacts to solid sulfur cyclooctasulfur; the molecule shape is similar to that of ozone. S3 is found in sulfur vapour, comprising 10 % of vapour species at 1,333 Pa.. It is cherry red in colour, with a bent structure, similar to ozone, O3; the bonds between the atoms are not full double bonds, the molecule can be thought of as a resonance between two states, in each of which one of the end atoms has a negative formal charge while the central atom has a positive formal charge, making it valence isoelectronic to ozone. The molecule has a distance between sulfur atoms of 191.70 ±.01 pm and angle at the central atom of 117.36°±0.006°. However, cyclic S3, where the sulfur atoms are arranged in an equilateral triangle with three bonds, should in theory be lower in energy than the bent structure observed.
The name thiozone was invented by Hugo Erdmann in 1908 who hypothesized that S3 made up a large proportion of liquid sulfur. However its existence was unproven until the experiments of J. Berkowitz in 1964. Using mass spectrometry, he showed. Above 1,200 °C S3 is the second most common molecule after S2 in sulfur gas. In liquid sulfur the molecule is not common until the temperature is high, such as 500 °C. However, small molecules like this contribute to most of the reactivity of liquid sulfur. S3 has an absorption peak of 425 nm with a tail extending into blue light. S3 can be made by photolysis of S3Cl2 embedded in a glass or matrix of solid noble gas, it occurs on Io in volcanic emissions. S3 is to appear in the atmosphere of Venus at heights of 20 to 30 km, where it is in thermal equilibrium with S2 and S4; the reddish colour of Venus atmosphere at lower levels is to be due to S3. S3 reacts with carbon monoxide CO to make carbon oxysulfide and S2. Tungsten and group-8 metal carbonyls, for example iron carbonyl, in theory can replace a carbonyl group with S3.
Formation of compounds with a defined number of sulfur atoms is possible: S3 + S−2 → S−5 S3 can form the radical anion S−3, which has an intense blue colour. The ion is called thiozonide, by analogy with the ozonide anion, O−3; the gemstone lapis lazuli and the mineral lazurite contain S−3. International Klein Blue, developed by Yves Klein contains the S−3 radical anion; this is valence isoelectronic with the ozonide ion. The spectrum of the colour shows a strong absorption band at 610–620 nm or 2.07 eV. The blue colour is due to the C2A2 transition to the X2B1 electronic state in the ion; the Raman frequency is 523 cm−1 and another infrared absorption is at 580 cm−1. The S−3 ion has been shown to be stable in aqueous solution under pressure of 0.5 GPa, is expected to occur at depth in the earth's crust where subduction or high pressure metamorphism occurs. This ion is important in movement of copper and gold in hydrothermal fluids. Lithium hexasulfide with tetramethylenediamine solvation reacts with acetone or donor solvents to form S−3.
The S−3 radical anion was made by reducing sulfur gas with Zn2+ in a matrix. The material is blue coloured when dry and changes colour to green and yellow in the presence of trace amounts of water. Another way to make it is with polysulfide dissolved in hexamethylphosphoramide where it gives a blue colour. Other methods of production of S−3 include reacting sulfur with dampened magnesium oxide. Raman spectroscopy can be used to identify S−3, it can be used non-destructively in paintings; the bands are 549 cm−1 for symmetric stretch, 585 cm−1 for asymmetric stretch, 259 cm−1 for bending. Natural materials can contain S−2 which has an optical absorption at 390 nm and Raman band at 590 cm−1; the trisulfide ion, S2−3 is part of the polysulfide series. The sulfur chain is bent at an angle of 107° 53'. SrS3 has a bond length in the trisulfide ion of 0.205 nm. The bonds are single, it is isoelectronic to sulfur dichloride
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
Allotropy or allotropism is the property of some chemical elements to exist in two or more different forms, in the same physical state, known as allotropes of the elements. Allotropes are different structural modifications of an element. For example, the allotropes of carbon include diamond, graphite and fullerenes; the term allotropy is used for elements only, not for compounds. The more general term, used for any crystalline material, is polymorphism. Allotropy refers only to different forms of an element within the same phase. For some elements, allotropes have different molecular formulae despite difference in phase. Other elements do not maintain distinct allotropes in different phases; the concept of allotropy was proposed in 1841 by the Swedish scientist Baron Jöns Jakob Berzelius. The term is derived from Greek, Modern άλλοτροπἱα, meaning'variability, changeableness'. After the acceptance of Avogadro's hypothesis in 1860, it was understood that elements could exist as polyatomic molecules, two allotropes of oxygen were recognized as O2 and O3.
In the early 20th century, it was recognized that other cases such as carbon were due to differences in crystal structure. By 1912, Ostwald noted that the allotropy of elements is just a special case of the phenomenon of polymorphism known for compounds, proposed that the terms allotrope and allotropy be abandoned and replaced by polymorph and polymorphism. Although many other chemists have repeated this advice, IUPAC and most chemistry texts still favour the usage of allotrope and allotropy for elements only. Allotropes are different structural forms of the same element and can exhibit quite different physical properties and chemical behaviours; the change between allotropic forms is triggered by the same forces that affect other structures, i.e. pressure and temperature. Therefore, the stability of the particular allotropes depends on particular conditions. For instance, iron changes from a body-centered cubic structure to a face-centered cubic structure above 906 °C, tin undergoes a modification known as tin pest from a metallic form to a semiconductor form below 13.2 °C.
As an example of allotropes having different chemical behaviour, ozone is a much stronger oxidizing agent than dioxygen. Elements capable of variable coordination number and/or oxidation states tend to exhibit greater numbers of allotropic forms. Another contributing factor is the ability of an element to catenate. Examples of allotropes include: Among the metallic elements that occur in nature in significant quantities half are allotropic at ambient pressure: Li, Be, Na, Ca, Ti, Mn, Fe, Co, Sr, Y, Zr, Sn, La, Ce, Pr, Nd, Sm, Gd, Tb, Dy, Yb, Hf, Tl, Th, Pa and U; some phase transitions between allotropic forms of technologically relevant metals are those of Ti at 882 °C, Fe at 912 °C and 1394 °C, Co at 422 °C, Zr at 863 °C, Sn at 13 °C and U at 668 °C and 776 °C. Cerium, samarium and ytterbium have three allotropes. Praseodymium, neodymium and terbium have two allotropes. Plutonium has six distinct solid allotropes under "normal" pressures, their densities vary within a ratio of some 4:3, which vastly complicates all kinds of work with the metal.
A seventh plutonium allotrope exists at high pressures. The transuranium metals Np, Am, Cm are allotropic. Promethium, americium and californium have three allotropes each. In 2017, the concept of nanoallotropy was proposed by Prof. Rafal Klajn of the Organic Chemistry Department of the Weizmann Institute of Science. Nanoallotropes, or allotropes of nanomaterials, are nanoporous materials that have the same chemical composition, but differ in their architecture at the nanoscale; such nanoallotropes may help create ultra-small electronic devices and find other industrial applications. The different nanoscale architectures translate into different properties, as was demonstrated for surface-enhanced Raman scattering performed on several different nanoallotropes of gold. A two-step method for generating nanoallotropes was created. Isomer Polymorphism Superdense carbon allotropes Chisholm, Hugh, ed.. "Allotropy". Encyclopædia Britannica. Cambridge University Press. "Untitled". Archived from the original on January 31, 2008.
Retrieved January 6, 2017. CS1 maint: BOT: original-url status unknown Allotropes – Chemistry Encyclopedia
Hydrogen is a chemical element with symbol H and atomic number 1. With a standard atomic weight of 1.008, hydrogen is the lightest element in the periodic table. Hydrogen is the most abundant chemical substance in the Universe, constituting 75% of all baryonic mass. Non-remnant stars are composed of hydrogen in the plasma state; the most common isotope of hydrogen, termed protium, has no neutrons. The universal emergence of atomic hydrogen first occurred during the recombination epoch. At standard temperature and pressure, hydrogen is a colorless, tasteless, non-toxic, nonmetallic combustible diatomic gas with the molecular formula H2. Since hydrogen forms covalent compounds with most nonmetallic elements, most of the hydrogen on Earth exists in molecular forms such as water or organic compounds. Hydrogen plays a important role in acid–base reactions because most acid-base reactions involve the exchange of protons between soluble molecules. In ionic compounds, hydrogen can take the form of a negative charge when it is known as a hydride, or as a positively charged species denoted by the symbol H+.
The hydrogen cation is written as though composed of a bare proton, but in reality, hydrogen cations in ionic compounds are always more complex. As the only neutral atom for which the Schrödinger equation can be solved analytically, study of the energetics and bonding of the hydrogen atom has played a key role in the development of quantum mechanics. Hydrogen gas was first artificially produced in the early 16th century by the reaction of acids on metals. In 1766–81, Henry Cavendish was the first to recognize that hydrogen gas was a discrete substance, that it produces water when burned, the property for which it was named: in Greek, hydrogen means "water-former". Industrial production is from steam reforming natural gas, less from more energy-intensive methods such as the electrolysis of water. Most hydrogen is used near the site of its production, the two largest uses being fossil fuel processing and ammonia production for the fertilizer market. Hydrogen is a concern in metallurgy as it can embrittle many metals, complicating the design of pipelines and storage tanks.
Hydrogen gas is flammable and will burn in air at a wide range of concentrations between 4% and 75% by volume. The enthalpy of combustion is −286 kJ/mol: 2 H2 + O2 → 2 H2O + 572 kJ Hydrogen gas forms explosive mixtures with air in concentrations from 4–74% and with chlorine at 5–95%; the explosive reactions may be triggered by heat, or sunlight. The hydrogen autoignition temperature, the temperature of spontaneous ignition in air, is 500 °C. Pure hydrogen-oxygen flames emit ultraviolet light and with high oxygen mix are nearly invisible to the naked eye, as illustrated by the faint plume of the Space Shuttle Main Engine, compared to the visible plume of a Space Shuttle Solid Rocket Booster, which uses an ammonium perchlorate composite; the detection of a burning hydrogen leak may require a flame detector. Hydrogen flames in other conditions are blue; the destruction of the Hindenburg airship was a notorious example of hydrogen combustion and the cause is still debated. The visible orange flames in that incident were the result of a rich mixture of hydrogen to oxygen combined with carbon compounds from the airship skin.
H2 reacts with every oxidizing element. Hydrogen can react spontaneously and violently at room temperature with chlorine and fluorine to form the corresponding hydrogen halides, hydrogen chloride and hydrogen fluoride, which are potentially dangerous acids; the ground state energy level of the electron in a hydrogen atom is −13.6 eV, equivalent to an ultraviolet photon of 91 nm wavelength. The energy levels of hydrogen can be calculated accurately using the Bohr model of the atom, which conceptualizes the electron as "orbiting" the proton in analogy to the Earth's orbit of the Sun. However, the atomic electron and proton are held together by electromagnetic force, while planets and celestial objects are held by gravity; because of the discretization of angular momentum postulated in early quantum mechanics by Bohr, the electron in the Bohr model can only occupy certain allowed distances from the proton, therefore only certain allowed energies. A more accurate description of the hydrogen atom comes from a purely quantum mechanical treatment that uses the Schrödinger equation, Dirac equation or the Feynman path integral formulation to calculate the probability density of the electron around the proton.
The most complicated treatments allow for the small effects of special relativity and vacuum polarization. In the quantum mechanical treatment, the electron in a ground state hydrogen atom has no angular momentum at all—illustrating how the "planetary orbit" differs from electron motion. There exist two different spin isomers of hydrogen diatomic molecules that differ by the relative spin of their nuclei. In the orthohydrogen form, the spins of the two protons are parallel and form a triplet state with a molecular spin quantum number of 1. At standard temperature and pressure, hydrogen gas contains about 25% of the para form and 75% of the ortho form known as the "normal form"; the equilibrium ratio of orthohydrogen to parahydrogen depends on temperature, but because the ortho form is an excited state and has a higher energy
Oxygen is the chemical element with the symbol O and atomic number 8. It is a member of the chalcogen group on the periodic table, a reactive nonmetal, an oxidizing agent that forms oxides with most elements as well as with other compounds. By mass, oxygen is the third-most abundant element in the universe, after helium. At standard temperature and pressure, two atoms of the element bind to form dioxygen, a colorless and odorless diatomic gas with the formula O2. Diatomic oxygen gas constitutes 20.8% of the Earth's atmosphere. As compounds including oxides, the element makes up half of the Earth's crust. Dioxygen is used in cellular respiration and many major classes of organic molecules in living organisms contain oxygen, such as proteins, nucleic acids and fats, as do the major constituent inorganic compounds of animal shells and bone. Most of the mass of living organisms is oxygen as a component of water, the major constituent of lifeforms. Oxygen is continuously replenished in Earth's atmosphere by photosynthesis, which uses the energy of sunlight to produce oxygen from water and carbon dioxide.
Oxygen is too chemically reactive to remain a free element in air without being continuously replenished by the photosynthetic action of living organisms. Another form of oxygen, ozone absorbs ultraviolet UVB radiation and the high-altitude ozone layer helps protect the biosphere from ultraviolet radiation. However, ozone present at the surface is a byproduct of thus a pollutant. Oxygen was isolated by Michael Sendivogius before 1604, but it is believed that the element was discovered independently by Carl Wilhelm Scheele, in Uppsala, in 1773 or earlier, Joseph Priestley in Wiltshire, in 1774. Priority is given for Priestley because his work was published first. Priestley, called oxygen "dephlogisticated air", did not recognize it as a chemical element; the name oxygen was coined in 1777 by Antoine Lavoisier, who first recognized oxygen as a chemical element and characterized the role it plays in combustion. Common uses of oxygen include production of steel and textiles, brazing and cutting of steels and other metals, rocket propellant, oxygen therapy, life support systems in aircraft, submarines and diving.
One of the first known experiments on the relationship between combustion and air was conducted by the 2nd century BCE Greek writer on mechanics, Philo of Byzantium. In his work Pneumatica, Philo observed that inverting a vessel over a burning candle and surrounding the vessel's neck with water resulted in some water rising into the neck. Philo incorrectly surmised that parts of the air in the vessel were converted into the classical element fire and thus were able to escape through pores in the glass. Many centuries Leonardo da Vinci built on Philo's work by observing that a portion of air is consumed during combustion and respiration. In the late 17th century, Robert Boyle proved. English chemist John Mayow refined this work by showing that fire requires only a part of air that he called spiritus nitroaereus. In one experiment, he found that placing either a mouse or a lit candle in a closed container over water caused the water to rise and replace one-fourteenth of the air's volume before extinguishing the subjects.
From this he surmised that nitroaereus is consumed in both combustion. Mayow observed that antimony increased in weight when heated, inferred that the nitroaereus must have combined with it, he thought that the lungs separate nitroaereus from air and pass it into the blood and that animal heat and muscle movement result from the reaction of nitroaereus with certain substances in the body. Accounts of these and other experiments and ideas were published in 1668 in his work Tractatus duo in the tract "De respiratione". Robert Hooke, Ole Borch, Mikhail Lomonosov, Pierre Bayen all produced oxygen in experiments in the 17th and the 18th century but none of them recognized it as a chemical element; this may have been in part due to the prevalence of the philosophy of combustion and corrosion called the phlogiston theory, the favored explanation of those processes. Established in 1667 by the German alchemist J. J. Becher, modified by the chemist Georg Ernst Stahl by 1731, phlogiston theory stated that all combustible materials were made of two parts.
One part, called phlogiston, was given off when the substance containing it was burned, while the dephlogisticated part was thought to be its true form, or calx. Combustible materials that leave little residue, such as wood or coal, were thought to be made of phlogiston. Air did not play a role in phlogiston theory, nor were any initial quantitative experiments conducted to test the idea. Polish alchemist and physician Michael Sendivogius in his work De Lapide Philosophorum Tractatus duodecim e naturae fonte et manuali experientia depromti described a substance contained in air, referring to it as'cibus vitae', this substance is identical with oxygen. Sendivogius, during his experiments performed between 1598 and 1604, properly recognized that the substance is equivalent to the gaseous byproduct released by the thermal decomposition of potassium nitrate. In Bugaj’s view, the isolation of oxygen and the proper association of the substance to that part of air, required for life, lends sufficient weight to the discovery of oxygen by Sendivogius.
Diatomic molecules are molecules composed of only two atoms, of the same or different chemical elements. The prefix di- is of Greek origin, meaning "two". If a diatomic molecule consists of two atoms of the same element, such as hydrogen or oxygen it is said to be homonuclear. Otherwise, if a diatomic molecule consists of two different atoms, such as carbon monoxide or nitric oxide, the molecule is said to be heteronuclear; the only chemical elements that form stable homonuclear diatomic molecules at standard temperature and pressure are the gases hydrogen, oxygen and chlorine. The noble gases are gases at STP, but they are monatomic; the homonuclear diatomic gases and noble gases together are called "elemental gases" or "molecular gases", to distinguish them from other gases that are chemical compounds. At elevated temperatures, the halogens bromine and iodine form diatomic gases. All halogens have been observed as diatomic molecules, except for astatine, uncertain; the mnemonics BrINClHOF, pronounced "Brinklehof", HONClBrIF, pronounced "Honkelbrif", HOFBrINCl have been coined to aid recall of the list of diatomic elements.
Other elements form diatomic molecules when evaporated, but these diatomic species repolymerize when cooled. Heating elemental phosphorus gives diphosphorus, P2. Sulfur vapor is disulfur. Dilithium is known in the gas phase. Ditungsten and dimolybdenum form with sextuple bonds in the gas phase; the bond in a homonuclear diatomic molecule is non-polar. Dirubidium is diatomic. All other diatomic molecules are chemical compounds of two different elements. Many elements can combine to form heteronuclear diatomic molecules, depending on temperature and pressure; some examples include, gases carbon monoxide, nitric oxide, hydrogen chloride. Many 1:1 binary compounds are not considered diatomic because they are polymeric at room temperature, but they form diatomic molecules when evaporated, for example gaseous MgO, SiO, many others. Hundreds of diatomic molecules have been identified in the environment of the Earth, in the laboratory, in interstellar space. About 99% of the Earth's atmosphere is composed of two species of diatomic molecules: nitrogen and oxygen.
The natural abundance of hydrogen in the Earth's atmosphere is only of the order of parts per million, but H2 is the most abundant diatomic molecule in the universe. The interstellar medium is, dominated by hydrogen atoms. Diatomic elements played an important role in the elucidation of the concepts of element and molecule in the 19th century, because some of the most common elements, such as hydrogen and nitrogen, occur as diatomic molecules. John Dalton's original atomic hypothesis assumed that all elements were monatomic and that the atoms in compounds would have the simplest atomic ratios with respect to one another. For example, Dalton assumed water's formula to be HO, giving the atomic weight of oxygen as eight times that of hydrogen, instead of the modern value of about 16; as a consequence, confusion existed regarding atomic weights and molecular formulas for about half a century. As early as 1805, Gay-Lussac and von Humboldt showed that water is formed of two volumes of hydrogen and one volume of oxygen, by 1811 Amedeo Avogadro had arrived at the correct interpretation of water's composition, based on what is now called Avogadro's law and the assumption of diatomic elemental molecules.
However, these results were ignored until 1860 due to the belief that atoms of one element would have no chemical affinity toward atoms of the same element, partly due to apparent exceptions to Avogadro's law that were not explained until in terms of dissociating molecules. At the 1860 Karlsruhe Congress on atomic weights, Cannizzaro resurrected Avogadro's ideas and used them to produce a consistent table of atomic weights, which agree with modern values; these weights were an important prerequisite for the discovery of the periodic law by Dmitri Mendeleev and Lothar Meyer. Diatomic molecules are in their lowest or ground state, which conventionally is known as the X state; when a gas of diatomic molecules is bombarded by energetic electrons, some of the molecules may be excited to higher electronic states, as occurs, for example, in the natural aurora. Such excitation can occur when the gas absorbs light or other electromagnetic radiation; the excited states are unstable and relax back to the ground state.
Over various short time scales after the excitation, transitions occur from higher to lower electronic states and to the ground state, in each transition results a photon is emitted. This emission is known as fluorescence. Successively higher electronic states are conventionally named A, B, C, etc.. The excitation energy must be greater than or equal to the energy of the electronic state in order for the excitation to occur. In quantum theory, an electronic state of a diatomic molecule is represented by 2 S + 1 Λ ( v