In chemistry, polarity is a separation of electric charge leading to a molecule or its chemical groups having an electric dipole moment, with a negatively charged end and a positively charged end. Polar molecules must contain polar bonds due to a difference in electronegativity between the bonded atoms. A polar molecule with two or more polar bonds must have a geometry, asymmetric in at least one direction, so that the bond dipoles do not cancel each other. Polar molecules interact through dipole–dipole intermolecular forces and hydrogen bonds. Polarity underlies a number of physical properties including surface tension and melting and boiling points. Not all atoms attract electrons with the same force; the amount of "pull" an atom exerts on its electrons is called its electronegativity. Atoms with high electronegativities – such as fluorine and nitrogen – exert a greater pull on electrons than atoms with lower electronegativities such as alkali metals and alkaline earth metals. In a bond, this leads to unequal sharing of electrons between the atoms, as electrons will be drawn closer to the atom with the higher electronegativity.
Because electrons have a negative charge, the unequal sharing of electrons within a bond leads to the formation of an electric dipole: a separation of positive and negative electric charge. Because the amount of charge separated in such dipoles is smaller than a fundamental charge, they are called partial charges, denoted as δ+ and δ−; these symbols were introduced by Sir Christopher Ingold and Dr. Edith Hilda Ingold in 1926; the bond dipole moment is calculated by multiplying the amount of charge separated and the distance between the charges. These dipoles within molecules can interact with dipoles in other molecules, creating dipole-dipole intermolecular forces. Bonds can fall between one of two extremes – being nonpolar or polar. A nonpolar bond occurs when the electronegativities are identical and therefore possess a difference of zero. A polar bond is more called an ionic bond, occurs when the difference between electronegativities is large enough that one atom takes an electron from the other.
The terms "polar" and "nonpolar" are applied to covalent bonds, that is, bonds where the polarity is not complete. To determine the polarity of a covalent bond using numerical means, the difference between the electronegativity of the atoms is used. Bond polarity is divided into three groups that are loosely based on the difference in electronegativity between the two bonded atoms. According to the Pauling scale: Nonpolar bonds occur when the difference in electronegativity between the two atoms is less than 0.5 Polar bonds occur when the difference in electronegativity between the two atoms is between 0.5 and 2.0 Ionic bonds occur when the difference in electronegativity between the two atoms is greater than 2.0Pauling based this classification scheme on the partial ionic character of a bond, an approximate function of the difference in electronegativity between the two bonded atoms. He estimated that a difference of 1.7 corresponds to 50% ionic character, so that a greater difference corresponds to a bond, predominantly ionic.
As a quantum-mechanical description, Pauling proposed that the wave function for a polar molecule AB is a linear combination of wave functions for covalent and ionic molecules: ψ = aψ + bψ. The amount of covalent and ionic character depends on the values of the squared coefficients a2 and b2. While the molecules can be described as "polar covalent", "nonpolar covalent", or "ionic", this is a relative term, with one molecule being more polar or more nonpolar than another. However, the following properties are typical of such molecules. A molecule is composed of one or more chemical bonds between molecular orbitals of different atoms. A molecule may be polar either as a result of polar bonds due to differences in electronegativity as described above, or as a result of an asymmetric arrangement of nonpolar covalent bonds and non-bonding pairs of electrons known as a full molecular orbital. A polar molecule has a net dipole as a result of the opposing charges from polar bonds arranged asymmetrically.
Water is an example of a polar molecule since it has a slight positive charge on one side and a slight negative charge on the other. The dipoles do not cancel out resulting in a net dipole. Due to the polar nature of the water molecule itself, polar molecules are able to dissolve in water. Other examples include sugars, which have many polar oxygen–hydrogen groups and are overall polar. If the bond dipole moments of the molecule do not cancel, the molecule is polar. For example, the water molecule contains two polar O−H bonds in a bent geometry; the bond dipole moments do not cancel, so that the molecule forms a molecular dipole with its negative pole at the oxygen and its positive pole midway between the two hydrogen atoms. In the figure each bond joins the central O atom with a negative charge to an H atom with a positive charge; the hydrogen fluoride, HF, molecule is polar by virtue of polar covalent bonds – in the covalent bond electrons are displaced toward the more electronegative fluorine atom.
Ammonia, NH3, molecule. The molecule has two lone electrons in an orbital, that points towards the fourth apex of the approximate tetrahedron; this orbital is not participating in covalent bonding.
The angstrom or ångström is a unit of length equal to 10−10 m. Its symbol is a letter of the Swedish alphabet; the angstrom is not a part of the SI system of units, but it can be considered part of the metric system. While deprecated by the IBWM and the NIST, the unit is still used in the natural sciences and technology to express sizes of atoms, microscopic biological structures, lengths of chemical bonds, arrangement of atoms in crystals, wavelengths of electromagnetic radiation, dimensions of integrated circuit parts; the atomic radii of phosphorus and chlorine are about 1 angstrom, while that of hydrogen is about 0.5 angstrom. Visible light has wavelengths in the range of 4000–7000 Å; the unit is named after the nineteenth-century Swedish physicist Anders Jonas Ångström. The IBWM and the NIST spell it as ångström; the symbol should always be "Å". The angstrom is used extensively in crystallography, solid-state physics and chemistry as a unit for d-spacings, cell parameters, inter-atomic distances and x-ray wavelengths, as these values are in the 1–10 Å range.
For example, the Inorganic Crystal Structure Database presents all these values using the angstrom. Anders Jonas Ångström was a pioneer in the field of spectroscopy, he is well known for his studies of astrophysics, heat transfer, terrestrial magnetism, the aurora borealis. In 1852, Ångström formulated in Optiska undersökningar, a law of absorption modified somewhat and known as Kirchhoff's law of thermal radiation. In 1868, Ångström created a chart of the spectrum of sunlight, in which he expressed the wavelengths of electromagnetic radiation in the electromagnetic spectrum in multiples of one ten-millionth of a millimetre Because the human eye is sensitive to wavelengths from about 4000 to 7000 Å, that choice of unit supported sufficiently accurate measurements of visible wavelengths without resorting to fractional numbers. Ångström's chart and table of wavelengths in the solar spectrum became used in solar physics, which adopted the unit and named it after him. It subsequently spread to the rest of astronomical spectroscopy, atomic spectroscopy, subsequently to other sciences that deal with atomic-scale structures.
Though intended to correspond to 10−10 metres, for precise spectral analysis, the angstrom had to be defined more than the metre, which until 1960 was still defined based on the length of a bar of metal held in Paris. The use of metal bars had been involved in an early error in the value of the angstrom of about one part in 6000. Ångström took the precaution of having the standard bar he used checked against a standard in Paris, but the metrologist Henri Tresca reported it to be so much shorter than it was that Ångström's corrected results were more in error than the uncorrected ones. In 1892–1895, Albert A. Michelson defined the angstrom so that the red line of cadmium was equal to 6438.47 angstroms. "In 1907, the International Union for Cooperation in Solar Research defined the international angstrom by declaring the wavelength of the red line of cadmium equal to 6438.4696 international angstroms, this definition was endorsed by the International Bureau of Weights and Measures in 1927. From 1927 to 1960, the angstrom remained a secondary unit of length for use in spectroscopy, defined separately from the metre.
In 1960, the metre itself was redefined in spectroscopic terms, which allowed the angstrom to be redefined as being 0.1 nanometres. The angstrom is internationally recognized, but is not a formal part of the International System of Units; the closest SI unit is the nanometre. The International Committee for Weights and Measures discourages its use, it is not included in the European Union's catalogue of units of measure that may be used within its internal market. For compatibility reasons, Unicode includes the formal symbol at U+212B Å ANGSTROM SIGN. However, the angstrom sign is normalized into U+00C5 Å LATIN CAPITAL LETTER A WITH RING ABOVE The Unicode consortium recommends to use the regular letter. Before digital typesetting, the angstrom was sometimes written as "A. U.". This use is evident in Bragg's paper on the structure of ice, which gives the c- and a-axis lattice constants as 4.52 A. U. and 7.34 A. U. respectively. Nowadays the atomic unit of length stands for bohrs, not angstroms. 100 picometres X unit Conversion of units
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
In chemistry and materials science the'coordination number' called ligancy, of a central atom in a molecule or crystal is the number of atoms, molecules or ions bonded to it. The ion/molecule/atom surrounding the central ion/molecule/atom is called a ligand; this number is determined somewhat differently for molecules than for crystals. For molecules and polyatomic ions the coordination number of an atom is determined by counting the other atoms to which it is bonded. For example, − has Cr3+ as its central cation, which has a coordination number of 6 and is described as hexacoordinate; however the solid-state structures of crystals have less defined bonds, in these cases a count of neighboring atoms is employed. The simplest method is one used in materials science; the usual value of the coordination number for a given structure refers to an atom in the interior of a crystal lattice with neighbors in all directions. In contexts where crystal surfaces are important, such as materials science and heterogeneous catalysis, the number of neighbors of an interior atom is the bulk coordination number, while the number of surface neighbors of an atom at the surface of the crystal is the surface coordination number.
In chemistry, coordination number, defined in 1893 by Alfred Werner, is the total number of neighbors of a central atom in a molecule or ion. Although a carbon atom has four chemical bonds in most stable molecules, the coordination number of each carbon is four in methane, three in ethylene, two in acetylene. In effect we count the first bond to each neighboring atom, but not the other bonds. In coordination complexes, only the first or sigma bond between each ligand and the central atom counts, but not any pi bonds to the same ligands. In tungsten hexacarbonyl, W6, the coordination number of tungsten is counted as six although pi as well as sigma bonding is important in such metal carbonyls; the most common coordination number for d-block transition metal complexes is 6, with an octahedral geometry. The observed range is 2 to 9. Metals in the f-block can accommodate higher coordination number due to their greater ionic radii and availability of more orbitals for bonding. Coordination numbers of 8 to 12 are observed for f-block elements.
For example, with bidentate nitrate ions as ligands, CeIV and ThIV form the 12-coordinate ions 2− and 2−. When the surrounding ligands are much smaller than the central atom higher coordination numbers may be possible. One computational chemistry study predicted a stable PbHe2+15 ion composed of a central lead ion coordinated with no fewer than 15 helium atoms. At the opposite extreme, steric shielding can give rise to unusually low coordination numbers. An rare instance of a metal adopting a coordination number of 1 occurs in the terphenyl-based arylthallium complex 2,6-Tipp2C6H3Tl, where Tipp is the 2,4,6-triisopropylphenyl group. For π-electron ligands such as the cyclopentadienide ion −, alkenes and the cyclooctatetraenide ion 2−, the number of atoms in the π-electron system that bind to the central atom is termed the hapticity. In ferrocene the hapticity, η, of each cyclopentadienide anion is five, Fe2. There are various ways of assigning the contribution made to the coordination number of the central iron atom by each cyclopentadienide ligand.
The contribution could be assigned as one since there is one ligand, or as five since there are five neighbouring atoms, or as three since there are three electron pairs involved. The count of electron pairs is taken. In order to determine the coordination number of an atom in a crystal, the crystal structure has first to be determined; this is achieved using neutron or electron diffraction. Once the positions of the atoms within the unit cell of the crystal are known the coordination number of an atom can be determined. For molecular solids or coordination complexes the units of the polyatomic species can be detected and a count of the bonds can be performed. Solids with lattice structures which includes metals and many inorganic solids can have regular structures where coordinating atoms are all at the same distance and they form the vertices of a coordination polyhedron. However, there are many such solids where the structures are irregular. In materials science, the bulk coordination number of a given atom in the interior of a crystal lattice is the number of nearest neighbours to a given atom.
For an atom at a surface of a crystal, the surface coordination number is always less than the bulk coordination number. The surface coordination number is dependent on the Miller indices of the surface. In a body-centered cubic crystal, the bulk coordination number is 8, for the surface, the surface coordination number is 4.α-Aluminium has a regular cubic close packed structure, where each aluminium atom has 12 nearest neighbors, 6 in the same plane and 3 above and below and the coordination polyhedron is a cuboctahedron. Α-Iron has a body centered cubic structure where each iron atom has 8 nearest neighbors situated at the corners of a cube. The two most common allotropes of carbon have different coordination numbers. In diamond, each carbon atom is at the centre of a regular tetrahedron formed by four other carbon atoms, the coordination number is four, as for methane. Graphite is made of two-dimensional layers in which each carbon is covalently bonded to three other carbons.
In molecular geometry, bond length or bond distance is the average distance between nuclei of two bonded atoms in a molecule. It is a transferable property of a bond between atoms of fixed types independent of the rest of the molecule. Bond length is related to bond order: when more electrons participate in bond formation the bond is shorter. Bond length is inversely related to bond strength and the bond dissociation energy: all other factors being equal, a stronger bond will be shorter. In a bond between two identical atoms, half the bond distance is equal to the covalent radius. Bond lengths are measured in the solid phase by means of X-ray diffraction, or approximated in the gas phase by microwave spectroscopy. A bond between a given pair of atoms may vary between different molecules. For example, the carbon to hydrogen bonds in methane are different from those in methyl chloride, it is however possible to make generalizations. A table with experimental single bonds for carbon to other elements is given below.
Bond lengths are given in picometers. By approximation the bond distance between two different atoms is the sum of the individual covalent radii; as a general trend, bond distances decrease across the row in the periodic table and increase down a group. This trend is identical to that of the atomic radius; the bond length between two atoms in a molecule depends not only on the atoms but on such factors as the orbital hybridization and the electronic and steric nature of the substituents. The carbon–carbon bond length in diamond is 154 pm, the largest bond length that exists for ordinary carbon covalent bonds. Since one atomic unit of length is 52.9177 pm, the C–C bond length is 2.91 atomic units, or three Bohr radii long. Unusually long bond lengths do exist. In one compound, tricyclobutabenzene, a bond length of 160 pm is reported; the current record holder is another cyclobutabenzene with length 174 pm based on X-ray crystallography. In this type of compound the cyclobutane ring would force 90° angles on the carbon atoms connected to the benzene ring where they ordinarily have angles of 120°.
The existence of a long C–C bond length of up to 290 pm is claimed in a dimer of two tetracyanoethylene dianions, although this concerns a 2-electron-4-center bond. This type of bonding has been observed in neutral phenalenyl dimers; the bond lengths of these so-called "pancake bonds" are up to 305 pm. Shorter than average C–C bond distances are possible: alkenes and alkynes have bond lengths of 133 and 120 pm due to increased s-character of the sigma bond. In benzene all bonds have the same length: 139 pm. Carbon–carbon single bonds increased s-character is notable in the central bond of diacetylene and that of a certain tetrahedrane dimer. In propionitrile the cyano group withdraws electrons resulting in a reduced bond length. Squeezing a C–C bond is possible by application of strain. An unusual organic compound exists called In-methylcyclophane with a short bond distance of 147 pm for the methyl group being squeezed between a triptycene and a phenyl group. In an in silico experiment a bond distance of 136 pm was estimated for neopentane locked up in fullerene.
The smallest theoretical C–C single bond obtained in this study is 131 pm for a hypothetical tetrahedrane derivative. The same study estimated that stretching or squeezing the C–C bond in an ethane molecule by 5 pm required 2.8 or 3.5 kJ/mol, respectively. Stretching or squeezing the same bond by 15 pm required an estimated 21.9 or 37.7 kJ/mol. Bond length tutorial
Linus Carl Pauling was an American chemist, peace activist, author and husband of American human rights activist Ava Helen Pauling. He published books, of which about 850 dealt with scientific topics. New Scientist called him one of the 20 greatest scientists of all time, as of 2000, he was rated the 16th most important scientist in history. Pauling was one of the founders of the fields of molecular biology, his contributions to the theory of the chemical bond include the concept of orbital hybridisation and the first accurate scale of electronegativities of the elements. Pauling worked on the structures of biological molecules, showed the importance of the alpha helix and beta sheet in protein secondary structure. Pauling's approach combined methods and results from X-ray crystallography, molecular model building and quantum chemistry, his discoveries inspired the work of James Watson, Francis Crick, Rosalind Franklin on the structure of DNA, which in turn made it possible for geneticists to crack the DNA code of all organisms.
In his years he promoted nuclear disarmament, as well as orthomolecular medicine, megavitamin therapy, dietary supplements. None of the latter have gained much acceptance in the mainstream scientific community. For his scientific work, Pauling was awarded the Nobel Prize in Chemistry in 1954. For his peace activism, he was awarded the Nobel Peace Prize in 1962, he is one of four individuals to have won more than one Nobel Prize. Of these, he is the only person to have been awarded two unshared Nobel Prizes, one of two people to be awarded Nobel Prizes in different fields, the other being Marie Curie. Pauling was born in Portland, the first-born child of Herman Henry William Pauling and Lucy Isabelle "Belle" Darling, he was named "Linus Carl", in honor of Lucy's father and Herman's father, Carl. In 1902, after his sister Pauline was born, Pauling's parents decided to move out of Portland, to find more affordable and spacious living quarters than their one-room apartment. Lucy stayed with her husband's parents in Lake Oswego until Herman brought the family to Salem, where he worked as a traveling salesman for the Skidmore Drug Company.
Within a year of Lucile's birth in 1904, Herman Pauling moved his family to Oswego, where he opened his own drugstore. He moved his family to Condon, Oregon, in 1905. By 1906, Herman Pauling was suffering from recurrent abdominal pain, he died of a perforated ulcer on June 11, 1910, leaving Lucy to care for Linus and Pauline. Pauling attributes his interest in becoming a chemist to being amazed by experiments conducted by a friend, Lloyd A. Jeffress, who had a small chemistry lab kit, he wrote: "I was entranced by chemical phenomena, by the reactions in which substances with strikingly different properties, appear. With an older friend, Lloyd Simon, Pauling set up Palmon Laboratories in Simon's basement, they approached local dairies offering to perform butterfat samplings at cheap prices but dairymen were wary of trusting two boys with the task, the business ended in failure. At age 15, the high school senior had enough credits to enter Oregon State University, known as Oregon Agricultural College.
Lacking two American history courses required for his high school diploma, Pauling asked the school principal if he could take the courses concurrently during the spring semester. Denied, he left Washington High School in June without a diploma; the school awarded him an honorary diploma 45 years after he was awarded two Nobel Prizes. Pauling held a number of jobs to earn money for his future college expenses, including working part-time at a grocery store for $8 per week, his mother arranged an interview with the owner of a number of manufacturing plants in Portland, Mr. Schwietzerhoff, who hired him as an apprentice machinist at a salary of $40 per month; this was soon raised to $50 per month. Pauling set up a photography laboratory with two friends. In September 1917, Pauling was admitted by Oregon State University, he resigned from the machinist's job and informed his mother, who saw no point in a university education, of his plans. In his first semester, Pauling registered for two courses in chemistry, two in mathematics, mechanical drawing, introduction to mining and use of explosives, modern English prose and military drill.
He founded the school's chapter of the Delta Upsilon fraternity. After his second year, he planned to take a job in Portland to help support his mother; the college offered him a position teaching quantitative analysis, a course he had just finished taking himself. He worked forty hours a week in the laboratory and classroom and earned $100 a month, enabling him to continue his studies. In his last two years at school, Pauling became aware of the work of Gilbert N. Lewis and Irving Langmuir on the electronic structure of atoms and their bonding to form molecules, he decided to focus his research on how the physical and chemical properties of substances are related to the structure of the atoms of which they are composed, becoming one of the founders of the new science of quantum chemistry. Engineering professor Samuel Graf selected Pauling to be his teaching assistant in a mechanics and materials course. During the winter of his senior year, Pauling taught a chemistry course for home economics majors.
It was in one of these classes that Pauling met his future wife
Oganesson is a synthetic chemical element with symbol Og and atomic number 118. It was first synthesized in 2002 at the Joint Institute for Nuclear Research in Dubna, near Moscow in Russia, by a joint team of Russian and American scientists. In December 2015, it was recognized as one of four new elements by the Joint Working Party of the international scientific bodies IUPAC and IUPAP, it was formally named on 28 November 2016. The name is in line with the tradition of honoring a scientist, in this case the nuclear physicist Yuri Oganessian, who has played a leading role in the discovery of the heaviest elements in the periodic table, it is one of only two elements named after a person, alive at the time of naming, the other being seaborgium. Oganesson has highest atomic mass of all known elements; the radioactive oganesson atom is unstable, since 2005, only five atoms of the nuclide 294Og have been detected. Although this allowed little experimental characterization of its properties and possible compounds, theoretical calculations have resulted in many predictions, including some surprising ones.
For example, although oganesson is a member of group 18 – the first synthetic element to be so – it may be reactive, unlike all the other elements of that group. It was thought to be a gas under normal conditions but is now predicted to be a solid due to relativistic effects. On the periodic table of the elements it is a p-block element and the last one of the period 7; the Danish physicist Niels Bohr was the first to consider the possibility of an element with an atomic number as high as 118, noting in 1922 that such an element would take its place in the periodic table below radon as the seventh noble gas. Following this, Aristid von Grosse wrote an article in 1965 predicting the properties of element 118; these were remarkably early predictions, given that it was not yet known how to produce elements artificially in 1922, that the existence of the island of stability had not yet been theorized in 1965. It was 80 years from Bohr's prediction before oganesson was synthesised, although its chemical properties have not been investigated to determine if it behaves as the heavier congener of radon.
In late 1998, Polish physicist Robert Smolańczuk published calculations on the fusion of atomic nuclei towards the synthesis of superheavy atoms, including oganesson. His calculations suggested that it might be possible to make oganesson by fusing lead with krypton under controlled conditions, that the fusion probability of that reaction would be close to the lead–chromium reaction that had produced element 106, seaborgium; this contradicted predictions that the cross-sections for reactions with lead or bismuth targets would go down exponentially as the atomic number of the resulting elements increased. In 1999, researchers at Lawrence Berkeley National Laboratory made use of these predictions and announced the discovery of livermorium and oganesson, in a paper published in Physical Review Letters, soon after the results were reported in Science; the researchers reported that they had performed the reaction 8636Kr + 20882Pb → 293118Og + n. The following year, they published a retraction after researchers at other laboratories were unable to duplicate the results and the Berkeley lab could not duplicate them either.
In June 2002, the director of the lab announced that the original claim of the discovery of these two elements had been based on data fabricated by principal author Victor Ninov. Newer experimental results and theoretical predictions have confirmed the exponential decrease in cross-sections with lead and bismuth targets as the atomic number of the resulting nuclide increases; the first genuine decay of atoms of oganesson was observed in 2002 at the Joint Institute for Nuclear Research in Dubna, Russia, by a joint team of Russian and American scientists. Headed by Yuri Oganessian, a Russian nuclear physicist of Armenian ethnicity, the team included American scientists of the Lawrence Livermore National Laboratory, California; the discovery was not announced because the decay energy of 294Og matched that of 212mPo, a common impurity produced in fusion reactions aimed at producing superheavy elements, thus announcement was delayed until after a 2005 confirmatory experiment aimed at producing more oganesson atoms.
On 9 October 2006, the researchers announced that they had indirectly detected a total of three nuclei of oganesson-294 produced via collisions of californium-249 atoms and calcium-48 ions. 24998Cf + 4820Ca → 294118Og + 3 n. In 2011, IUPAC evaluated the 2006 results of the Dubna–Livermore collaboration and concluded: "The three events reported for the Z = 118 isotope have good internal redundancy but with no anchor to known nuclei do not satisfy the criteria for discovery"; because of the small fusion reaction probability the experiment took four months and involved a beam dose of 2.5×1019 calcium ions that had to be shot at the californium target to produce the first recorded event believed to be the synthesis of oganesson. Researchers were confident that the results were not a false positive, since the chance that the detections were random events was estimated to be less than one part in 100000. In the experiments, the alpha-decay of three atoms of oganesson was observed. A fourth decay by direct spontaneous fission was proposed.
A half-life of 0.89 ms was calculated: 294Og decays into 290Lv by alpha decay. Since there were only three nuclei, the half-life derived from ob