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
Phosphorus mononitride is an inorganic compound with the chemical formula PN. Containing only phosphorus and nitrogen, this material is classified as a binary nitride, it is the first identified phosphorus compound in the interstellar medium. It is the atmospheres of Jupiter and Saturn. Triphosphorus pentanitride
Iron oxide or ferrous oxide is the inorganic compound with the formula FeO. Its mineral form is known as wüstite. One of several iron oxides, it is a black-colored powder, sometimes confused with rust, the latter of which consists of hydrated iron oxide. Iron oxide refers to a family of related non-stoichiometric compounds, which are iron deficient with compositions ranging from Fe0.84O to Fe0.95O. FeO can be prepared by the thermal decomposition of iron oxalate. FeC2O4 → FeO + CO2 + COThe procedure is conducted under an inert atmosphere to avoid the formation of ferric oxide. A similar procedure can be used for the synthesis of manganous oxide and stannous oxide. Stoichiometric FeO can be prepared by heating Fe0.95 O with metallic iron at 36 kbar. FeO is thermodynamically unstable below 575 °C, tending to disproportionate to metal and Fe3O4: 4FeO → Fe + Fe3O4 Iron oxide adopts the cubic, rock salt structure, where iron atoms are octahedrally coordinated by oxygen atoms and the oxygen atoms octahedrally coordinated by iron atoms.
The non-stoichiometry occurs because of the ease of oxidation of FeII to FeIII replacing a small portion of FeII with two thirds their number of FeIII, which take up tetrahedral positions in the close packed oxide lattice. Below 200 K there is a minor change to the structure which changes the symmetry to rhombohedral and samples become antiferromagnetic. Iron oxide makes up 9% of the Earth's mantle. Within the mantle, it may be electrically conductive, a possible explanation for perturbations in Earth's rotation not accounted for by accepted models of the mantle's properties. Iron dissolved in groundwater is in the reduced iron II form. If this groundwater comes in contact with oxygen at the surface, e.g. in natural springs, iron II is oxidised to iron III and forms insoluble hydroxides in water. Iron oxide is used as a pigment, it is FDA-approved for use in cosmetics and it is used in some tattoo inks. It can be used as a phosphate remover from home aquaria. Http://webmineral.com/data/Wustite.shtml
Methylidyne called carbyne, is an organic compound with the chemical formula CH•. Methylidyne is the simplest carbyne, it is a reactive gas, destroyed in ordinary conditions but is abundant in the interstellar medium. In October 2016, astronomers reported that the basic chemical ingredients of life – the carbon-hydrogen molecule, the carbon-hydrogen positive ion, the carbon ion – are the result of ultraviolet light from stars, rather than in other ways, such as the result of turbulent events related to supernovae and young stars, as thought earlier; these results have given new light to the formation of organic compounds in the early development of life on earth. The trivial name carbyne is the preferred IUPAC name; the systematic names methylidyne, hydridocarbon, valid IUPAC names, are constructed according to the substitutive and additive nomenclatures, respectively. Methylidyne is viewed as methane with three hydrogen atoms removed. By default, this name pays no regard to the radicality of methylidyne.
When the radicality is considered, the radical states with one unpaired electron are named methylylidene, whereas the radical excited states with three unpaired electrons are named methanetriyl. As an odd-electron species, CH is a radical; the ground state is a doublet. The first two excited states are a doublet; the quartet lies at 71 kJ above the ground state. Reactions of the doublet radical with non-radical species involves insertion or addition, whereas reactions of the quartet radical involves only abstraction. • + H2O → • + H2 or • 3• + H2O → + • Methylidyne-like species are implied intermediates in the Fischer–Tropsch process, the hydrogenation of CO into hydrocarbons. Methylidyne entities are assumed to bond to the catalyst's surface. A hypothetical sequence is: MnCO + 1/2 H2 → MnCOH MnCOH + H2 → MnCH + H2OA molecular example of an MnCH is HCCo39. MnCH + 1/2 H2 → MnCH2The methylene ligand is poised couple to CO or to another methylene, thereby growing the C–C chain; the methylylidyne group can exhibit both Lewis basic character.
Such behavior is only of theoretical interest. Methylidine can be prepared from bromoform. Methylene group Methylene bridge
Norman Neill Greenwood FRS CChem FRSC was an Australian-British chemist and Emeritus Professor at the University of Leeds. He is best known for the innovative textbook Chemistry of the Elements, co-authored with Alan Earnshaw, first published in 1984. After attending University High School, Greenwood read Chemistry at the University of Melbourne and graduated with a BSc in 1945 and an MSc in 1948. In 1948, he was awarded the Exhibition of 1851 Scholarship to enable him to read for a PhD at Sidney Sussex College, Cambridge under the supervision of Harry Julius Emeléus, he received the PhD in 1951. Greenwood was a senior research fellow at the Atomic Energy Research Establishment from 1951 until 1953 when he was appointed a lecturer at the University of Nottingham, his first PhD student at Nottingham was Kenneth Wade. Professor William Wynne-Jones, the Chairman of the School of Chemistry at Kings College, recruited Greenwood to the first established chair of inorganic chemistry in the country in 1961.
Greenwood was appointed professor and head of the Department of Inorganic and Structural Chemistry at the University of Leeds in 1971, a post which he held until his retirement in 1990 when he was given the title emeritus professor. Greenwood was elected a fellow of the Royal Society in 1987, his wide-ranging researches in inorganic and structural chemistry have made major advances in the chemistry of boron hydrides and other main-group element compounds. He pioneered the application of Mössbauer spectroscopy to problems in chemistry, he was a prolific writer and inspirational lecturer on chemical and educational themes, has held numerous visiting professorships throughout the world. He was appointed by NASA as principal investigator in the study of lunar rocks, he served as chairman of the IUPAC Commission on Atomic Weights from 1970 to 1975 and as president of the IUPAC Inorganic Chemistry Division. Greenwood, N. N.. Principles of Atomic Orbitals – Monograph for Teachers. Royal Society of Chemistry.
P. 48. ISBN 9780854040285. Greenwood, N. N.. Ionic crystals, lattice defects and nonstoichiometry. Butterworths. P. 194. Greenwood, N. N. C.. Mössbauer Spectroscopy. Chapman and Hall. P. 659. Greenwood, Norman N.. Chemistry of the Elements. Butterworth-Heinemann. P. 1340. ISBN 978-0-08-037941-8. Greenwood, N. N.. Recollections of a Scientist Volume 1. Boyhood and Youth in Australia. Xlibris Corporation. P. 288. ISBN 1-4691-7935-0. Greenwood, N. N.. Recollections of a Scientist, Volume 2: Expanding Horizons: England and Europe. Xlibris Corporation. P. 438. ISBN 978-1477151860. Editor: Spectroscopic Properties of Inorganic and Organometallic Compounds, Royal Society of Chemistry, Volume 1 to Volume 9 Norman Greenwood tells his life story at Web of Stories
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
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