A nitrogen atom has seven electrons. In the ground state, they are arranged in the electron configuration 1s2s2p2p2p. It therefore has five valence electrons in the 2s and 2p orbitals, three of which (the p-electrons) are unpaired. It has one of the highest electronegativities among the elements (3.04 on the Pauling scale), exceeded only by chlorine (3.16), oxygen (3.44), and fluorine (3.98). (The light noble gases, helium, neon, and argon, would presumably also be more electronegative, and in fact are on the Allen scale.) Following periodic trends, its single-bond covalent radius of 71 pm is smaller than those of boron (84 pm) and carbon (76 pm), while it is larger than those of oxygen (66 pm) and fluorine (57 pm).
Atomic
Nitrogen
Nitrogen | Atomic
It has seven electrons and one of the highest electronegativities among the elements. It is a colourless and odorless diatomic gas. It is the lightest member of group 15 of the periodic table and is a common element in the universe. Colourless and odorless
the general surface wind circulation associated with the high-pressure center (h) is most likely
user
A high-pressure area, high ,or anticyclone.
assistant
how is it formed?
user
Winds within high-pressure areas flow outward from the higher pressure areas near their centers towards the lower pressure areas further from their centers. Gravity adds to the forces. The winds are further twisted due to Coriolis effect.
assistant
could i know more about the effect mentioned above?
user
This is an apparent, but fictitious, force that arises when an observer is in a rotating reference frame, due to the conservation of angular momentum of the air as it moves towards or away from the Earth's axis of rotation.
assistant
how is it related to the wind conditions?
user
They are associated with light winds at the surface and subsidence through the lower portion of the troposphere.
assistant
what is the last term used above?
user
It is the lowest layer of earth's atmosphere.
assistant
what does it consist of?
user
Nitrogrn, oxygen, argon, carbon dioxide and small amounts of other gases.
assistant
who discovered the former from the list above?
user
Daniel Rutherford.
assistant
what are its properties?
user
The physical properties of N are quite different from other nuclei. Its properties along with several common nuclei are summarized in the below table.
Nitrogen-15 nuclear magnetic resonance spectroscopy | Physical properties
Atomic nitrogen, also known as active nitrogen, is highly reactive, being a triradical with three unpaired electrons. Free nitrogen atoms easily react with most elements to form nitrides, and even when two free nitrogen atoms collide to produce an excited N molecule, they may release so much energy on collision with even such stable molecules as carbon dioxide and water to cause homolytic fission into radicals such as CO and O or OH and H. Atomic nitrogen is prepared by passing an electric discharge through nitrogen gas at 0.1–2 mmHg, which produces atomic nitrogen along with a peach-yellow emission that fades slowly as an afterglow for several minutes even after the discharge terminates.
Nitrogen | Allotropes
Nitrogen is the chemical element with the symbol N and atomic number 7. It was first discovered and isolated by Scottish physician Daniel Rutherford in 1772. Although Carl Wilhelm Scheele and Henry Cavendish had independently done so at about the same time, Rutherford is generally accorded the credit because his work was published first. The name "nitrogène" was suggested by French chemist Jean-Antoine-Claude Chaptal in 1790 when it was found that nitrogen was present in nitric acid and nitrates. Antoine Lavoisier suggested instead the name "azote", from the Greek ἀζωτικός "no life", as it is an asphyxiant gas; this name is instead used in many languages, such as French, Italian, Russian, Romanian and Turkish, and appears in the English names of some nitrogen compounds such as hydrazine, azides and azo compounds.
Nitrogen | Introduction
It also results in very large electrostatic forces of attraction between the nucleus and the valence electrons in the 2s and 2p shells, resulting in very high electronegativities. Hypervalency is almost unknown in the 2p elements for the same reason, because the high electronegativity makes it difficult for a small nitrogen atom to be a central atom in an electron-rich three-center four-electron bond since it would tend to attract the electrons strongly to itself. Thus, despite nitrogen's position at the head of group 15 in the periodic table, its chemistry shows huge differences from that of its heavier congeners phosphorus, arsenic, antimony, and bismuth.
Nitrogen | Atomic
Nitrogen is the chemical element with atomic number 7, the symbol N and atomic mass 14.00674 u. Elemental nitrogen is a colorless, odorless, tasteless and mostly inert diatomic gas at standard conditions, constituting 78.08% by volume of Earth's atmosphere. The element nitrogen was discovered as a separable component of air, by Scottish physician Daniel Rutherford, in 1772. It occurs naturally in form of two isotopes: nitrogen-14 and nitrogen-15. Many industrially important compounds, such as ammonia, nitric acid, organic nitrates (propellants and explosives), and cyanides, contain nitrogen. The extremely strong bond in elemental nitrogen dominates nitrogen chemistry, causing difficulty for both organisms and industry in breaking the bond to convert the molecule into useful compounds, but at the same time causing release of large amounts of often useful energy when the compounds burn, explode, or decay back into nitrogen gas.
Period 2 element | Nitrogen
Rutherford discovered nitrogen by the isolation of the particle in 1772. When Joseph Black was studying the properties of carbon dioxide, he found that a candle would not burn in it. Black turned this problem over to his student at the time, Rutherford. Rutherford kept a mouse in a space with a confined quantity of air until it died. Then, he burned a candle in the remaining air until it went out. Afterwards, he burned phosphorus in that, until it would not burn. Then the air was passed through a carbon dioxide absorbing solution. The remaining component of the air did not support combustion, and a mouse could not live in it.
Daniel Rutherford | Isolation of nitrogen
In 1917 Rutherford bombarded nitrogen gas with alpha particles and observed hydrogen nuclei being emitted from the gas (Rutherford recognized these, because he had previously obtained them bombarding hydrogen with alpha particles, and observing hydrogen nuclei in the products). Rutherford concluded that the hydrogen nuclei emerged from the nuclei of the nitrogen atoms themselves (in effect, he had split a nitrogen). From his own work and the work of his students Bohr and Henry Moseley, Rutherford knew that the positive charge of any atom could always be equated to that of an integer number of hydrogen nuclei. This, coupled with the atomic mass of many elements being roughly equivalent to an integer number of hydrogen atoms - then assumed to be the lightest particles - led him to conclude that hydrogen nuclei were singular particles and a basic constituent of all atomic nuclei.
Atomic theory | Discovery of nuclear particles
In 1919–1920, Rutherford found that nitrogen and other light elements ejected a proton, which he called a "hydrogen atom", when hit with α (alpha) particles. This result showed Rutherford that hydrogen nuclei were a part of nitrogen nuclei (and by inference, probably other nuclei as well). Such a construction had been suspected for many years on the basis of atomic weights which were whole numbers of that of hydrogen; see Prout's hypothesis. Hydrogen was known to be the lightest element, and its nuclei presumably the lightest nuclei. Now, because of all these considerations, Rutherford decided that a hydrogen nucleus was possibly a fundamental building block of all nuclei, and also possibly a new fundamental particle as well, since nothing was known from the nucleus that was lighter.
Ernest Rutherford | Gold foil experiment
The nitride anion, N, is much larger at 146 pm, similar to that of the oxide (O: 140 pm) and fluoride (F: 133 pm) anions. The first three ionisation energies of nitrogen are 1.402, 2.856, and 4.577 MJ·mol, and the sum of the fourth and fifth is 16.920 MJ·mol. Due to these very high figures, nitrogen has no simple cationic chemistry. The lack of radial nodes in the 2p subshell is directly responsible for many of the anomalous properties of the first row of the p-block, especially in nitrogen, oxygen, and fluorine. The 2p subshell is very small and has a very similar radius to the 2s shell, facilitating orbital hybridisation.
Nitrogen | Atomic
A nitrogen atom has seven electrons. In the ground state, they are arranged in the electron configuration 1s2s2p2p2p. It therefore has five valence electrons in the 2s and 2p orbitals, three of which (the p-electrons) are unpaired. It has one of the highest electronegativities among the elements (3.04 on the Pauling scale), exceeded only by chlorine (3.16), oxygen (3.44), and fluorine (3.98). (The light noble gases, helium, neon, and argon, would presumably also be more electronegative, and in fact are on the Allen scale.) Following periodic trends, its single-bond covalent radius of 71 pm is smaller than those of boron (84 pm) and carbon (76 pm), while it is larger than those of oxygen (66 pm) and fluorine (57 pm).
Nitrogen | Atomic
Nitrogen may be usefully compared to its horizontal neighbours carbon and oxygen as well as its vertical neighbours in the pnictogen column, phosphorus, arsenic, antimony, and bismuth. Although each period 2 element from lithium to oxygen shows some similarities to the period 3 element in the next group (from magnesium to chlorine; these are known as diagonal relationships), their degree drops off abruptly past the boron–silicon pair. The similarities of nitrogen to sulfur are mostly limited to sulfur nitride ring compounds when both elements are the only ones present.
Nitrogen | Atomic
Nitrogen is a colourless, odourless, and relatively inert diatomic gas with a density of 1.251 × 10 g/cm (marginally heavier than air). It condenses to a colourless liquid at −195.795 °C and freezes into an ice- or snow-like solid at −210.00 °C. The solid form (density 0.85 g/cm; cf. lithium 0.534) has a hexagonal crystalline structure and is soft and easily crushed. Nitrogen is an insulator in all of its forms. It has a high ionisation energy (1402.3 kJ/mol), low electron affinity (–6.75 kJ/mol), and high electronegativity (3.04). The latter property manifests in the capacity of nitrogen to form usually strong hydrogen bonds, and its preference for forming complexes with metals having low electronegativities, small cationic radii, and often high charges (+3 or more).
Nonmetal | Group 15
After experimentation Rutherford traced the reaction to the nitrogen in air and found that when alpha particles were introduced into pure nitrogen gas, the effect was larger. In 1919 Rutherford assumed that the alpha particle knocked a proton out of nitrogen, turning it into carbon. After observing Blackett's cloud chamber images in 1925, Rutherford realized that the opposite was the case: after capture of the alpha particle, a proton is ejected, so that heavy oxygen, not carbon, is the end result i.e. Z is not decremented but incremented. This was the first reported nuclear reaction, N + α → O + p.
Proton | History
Throughout the 1920s, physicists assumed that the atomic nucleus was composed of protons and "nuclear electrons". Under this hypothesis, the nitrogen-14 (N) nucleus would be composed of 14 protons and 7 electrons, so that it would have a net charge of +7 elementary charge units and a mass of 14 atomic mass units. This nucleus would also be orbited by another 7 electrons, termed "external electrons" by Rutherford, to complete the N atom. However problems with the hypothesis soon became apparent. Ralph Kronig pointed out in 1926 that the observed hyperfine structure of atomic spectra was inconsistent with the proton–electron hypothesis.
Discovery of the neutron | Problems of the nuclear electrons hypothesis
Moseley, along with Charles Galton Darwin, Niels Bohr, and George de Hevesy, proposed that the nuclear charge ("Z") or atomic mass may be mathematically related to physical properties. The significance of these atomic properties was determined in the Geiger–Marsden experiment, in which the atomic nucleus and its charge were discovered.
History of the periodic table | Radioactivity, isotopes, and the Rutherford model
Based on the results of astronomical spectroscopy of nebula he proposed in 1911 the existence of several yet undiscovered elements. Coronium with an atomic weight of 0.51282, nebulium with a weight of 1.6281 and protofluorine with a weight of 2.361. Ira Sprague Bowen was able to attribute the spectroscopical lines of nebulium to doubly ionized oxygen making the new elements obsolete for their explanation.
John William Nicholson | Career
However, if the nitrogen-14 (N) nucleus was composed of 14 protons and 7 electrons, an odd number of spin-1/2 particles, then the resultant nuclear spin should be half-integer. Kronig therefore suggested that perhaps "protons and electrons do not retain their identity to the extent they do outside the nucleus".
Discovery of the neutron | Problems of the nuclear electrons hypothesis
Faster cooling may be obtained by plunging an object into a slush of liquid and solid nitrogen rather than liquid nitrogen alone.
Liquid nitrogen | Physical properties
At the beginning of the 20th century it was widely believed that atoms never change: an atom of one element stayed that way forever. Rutherford had already revealed the structure of an atom to consist mostly of empty space with a dense nucleus of protons at the centre, and Henry Mosley had shown that it is the number of protons that gives an atom its identity as a particular element. An atom of the element carbon has 6 protons, whilst an atom with 7 protons is one of nitrogen. Rutherford came to the conclusion that the number of protons in a radioactive element "could" change – through a process of decay where parts of the nucleus are ejected from the atom.
Chemistry: A Volatile History | Ernest Rutherford explains radioactivity
Indeed, the static dipole polarisability (α) of ununennium, a quantity for which the impacts of relativity are proportional to the square of the element's atomic number, has been calculated to be small and similar to that of sodium.
Ununennium | Atomic and physical
At atmospheric pressure, it sublimes at 56.5 °C. The solid state structure was determined by neutron diffraction at 77 K and 293 K.
Uranium hexafluoride | Physical properties
A study of Dalton's laboratory notebooks, discovered in the rooms of the Manchester Literary and Philosophical Society, concluded that so far from Dalton being led by his search for an explanation of the law of multiple proportions to the idea that chemical combination consists in the interaction of atoms of definite and characteristic weight, the idea of atoms arose in his mind as a purely physical concept, forced on him by study of the physical properties of the atmosphere and other gases. The first published indications of this idea are to be found at the end of his paper "On the Absorption of Gases by Water and other Liquids" already mentioned.
John Dalton | Atomic theory
The word "proton" is Greek for "first", and this name was given to the hydrogen nucleus by Ernest Rutherford in 1920. In previous years, Rutherford had discovered that the hydrogen nucleus (known to be the lightest nucleus) could be extracted from the nuclei of nitrogen by atomic collisions. Protons were therefore a candidate to be a fundamental particle, and hence a building block of nitrogen and all other heavier atomic nuclei.
Proton | Introduction
Daniel Rutherford (3 November 1749 – 15 December 1819) was a Scottish physician, chemist and botanist who is known for the isolation of nitrogen in 1772.
Daniel Rutherford | Introduction
This structure is caused by the influence of the nucleus on the dynamics of orbiting electrons. The magnetic moments of supposed "nuclear electrons" should produce hyperfine spectral line splittings similar to the Zeeman effect, but no such effects were observed. It seemed that the magnetic moment of the electron vanished when it was within the nucleus. While on a visit to Utrecht University in 1928, Kronig learned of a surprising aspect of the rotational spectrum of N. The precision measurement made by Leonard Ornstein, the director of Utrecht's Physical Laboratory, showed that the spin of nitrogen nucleus must be equal to one.
Discovery of the neutron | Problems of the nuclear electrons hypothesis
This proved eventually to be the case.
Atomic number | The Rutherford-Bohr model and van den Broek
Dalton published his first table of relative atomic weights containing six elements (hydrogen, oxygen, nitrogen, carbon, sulfur and phosphorus), relative to the weight of an atom of hydrogen conventionally taken as 1. Since these were only relative weights, they do not have a unit of weight attached to them. Dalton provided no indication in this paper how he had arrived at these numbers, but in his laboratory notebook, dated 6 September 1803, is a list in which he set out the relative weights of the atoms of a number of elements, derived from analysis of water, ammonia, carbon dioxide, etc.
John Dalton | Atomic weights
Rutherford called the gas (which we now know would have consisted primarily of nitrogen) "noxious air" or "phlogisticated air". Rutherford reported the experiment in 1772. He and Black were convinced of the validity of the phlogiston theory, so they explained their results in terms of it.
Daniel Rutherford | Isolation of nitrogen
They subsequently explained why the minor Mo and Mo isotopes of the element molybdenum had become dominant in supernovae stardust explaining an experimental riddle in stardust isotopic abundances.
Donald D. Clayton | Nuclear physics origin of the chemical elements (<a href="Nucleosynthesis">Nucleosynthesis</a>)
This is expected to be a similar configuration for , disproving a proposed tetrahedral (T) model in which all of the nitrogen atoms are equivalent.
Tetranitrogen | Characteristics
Although the converse is claimed to hold in Cohn's paper, this is known to be false. The term "atomic" is due to P. M. Cohn, who called an irreducible element of an integral domain an "atom".
Atomic domain | Introduction
The nitrogen compound sal ammoniac (ammonium chloride) has been known since the time of the Ancient Egyptians. In the 1760s two scientists, Henry Cavendish and Joseph Priestley, isolated nitrogen from air, but neither realized the presence of an undiscovered element. It was not until several years later, in 1772, that Daniel Rutherford realized that the gas was indeed nitrogen. The alchemist Hennig Brandt first discovered phosphorus in Hamburg in 1669. Brandt produced the element by heating evaporated urine and condensing the resulting phosphorus vapor in water. Brandt initially thought that he had discovered the Philosopher's Stone, but eventually realized that this was not the case.
Pnictogen | History
The story of the discovery of the neutron and its properties is central to the extraordinary developments in atomic physics that occurred in the first half of the 20th century, leading ultimately to the atomic bomb in 1945. In the 1911 Rutherford model, the atom consisted of a small positively charged massive nucleus surrounded by a much larger cloud of negatively charged electrons. In 1920, Rutherford suggested that the nucleus consisted of positive protons and neutrally-charged particles, suggested to be a proton and an electron bound in some way. Electrons were assumed to reside within the nucleus because it was known that beta radiation consisted of electrons emitted from the nucleus.
Neutron | Discovery
The nitrogen rule states that organic compounds containing exclusively hydrogen, carbon, nitrogen, oxygen, silicon, phosphorus, sulfur, and the halogens either have an odd nominal mass that indicates an odd number of nitrogen atoms are present or an even nominal mass that indicates an even number of nitrogen atoms are present in the molecular ion.
Mass (mass spectrometry) | Nitrogen rule
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John Dalton | Atomic theory
The nucleus was discovered in 1911, as a result of Ernest Rutherford's efforts to test Thomson's "plum pudding model" of the atom. The electron had already been discovered by J.J. Thomson himself. Knowing that atoms are electrically neutral, J.J.Thomson postulated that there must be a positive charge as well. In his plum pudding model, Thomson suggested that an atom consisted of negative electrons randomly scattered within a sphere of positive charge. Ernest Rutherford later devised an experiment with his research partner Hans Geiger and with help of Ernest Marsden, that involved the deflection of alpha particles (helium nuclei) directed at a thin sheet of metal foil.
Atomic nucleus | History
Cockcroft helped with the design and construction of helium liquefiers. In 1919, Rutherford had succeeded in disintegrating nitrogen atoms with alpha particles emitted from decaying radium atoms. This and subsequent experiments hinted at the structure of atomic nuclei. To explore it further, he needed an artificial means of creating particles with a velocity high enough to overcome the charge of the nucleus. This opened a new line of research at the Cavendish Laboratory. He assigned Cockcroft, Thomas Allibone and Ernest Walton to the problem. They built what became known as a Cockcroft–Walton accelerator. Mark Oliphant designed a proton source for them.
John Cockcroft | Nuclear research
In 1920 Rutherford gave a Bakerian lecture at the Royal Society entitled the "Nuclear Constitution of Atoms", a summary of recent experiments on atomic nuclei and conclusions as to the structure of atomic nuclei. By 1920, the existence of electrons within the atomic nucleus was widely assumed. It was assumed the nucleus consisted of hydrogen nuclei in number equal to the atomic mass. But since each hydrogen nucleus had charge +1, the nucleus required a smaller number of "internal electrons" each of charge -1 to give the nucleus its correct total charge. The mass of protons is about 1800 times greater than that of electrons, so the mass of the electrons is incidental in this computation.
Discovery of the neutron | Rutherford atom
Nitrogen is the lightest member of group 15 of the periodic table, often called the pnictogens. It is a common element in the universe, estimated at about seventh in total abundance in the Milky Way and the Solar System. At standard temperature and pressure, two atoms of the element bind to form dinitrogen, a colourless and odorless diatomic gas with the formula N. Dinitrogen forms about 78% of Earth's atmosphere, making it the most abundant uncombined element. Nitrogen occurs in all organisms, primarily in amino acids (and thus proteins), in the nucleic acids (DNA and RNA) and in the energy transfer molecule adenosine triphosphate.
Nitrogen | Introduction
The nitrogen rule states that organic compounds containing exclusively hydrogen, carbon, nitrogen, oxygen, silicon, phosphorus, sulfur, and the halogens either have 1) an "odd nominal mass" that indicates an "odd number" of nitrogen atoms are present or 2) an "even nominal mass" that indicates an "even number" of nitrogen atoms in the molecular formula of the molecular ion. The nitrogen rule is not a rule, "per se", as much as a general principle which may prove useful when attempting to solve organic mass spectrometry structures.
Nitrogen rule | Introduction
If Prout's hypothesis were true, something had to be neutralizing some of the charge of the hydrogen nuclei present in the nuclei of heavier atoms. In 1917, Rutherford succeeded in generating hydrogen nuclei from a nuclear reaction between alpha particles and nitrogen gas, and believed he had proven Prout's law. He called the new heavy nuclear particles protons in 1920 (alternate names being proutons and protyles). It had been immediately apparent from the work of Moseley that the nuclei of heavy atoms have more than twice as much mass as would be expected from their being made of hydrogen nuclei, and thus there was required a hypothesis for the neutralization of the extra protons presumed present in all heavy nuclei.
Atomic number | The proton and the idea of nuclear electrons
Molecular nitrogen is the gas form of nitrogen in the atmosphere. The chemical structure for nitrogen gas is N.
Ammoniacal nitrogen | Molecular nitrogen
Biologically mediated reactions (e.g., assimilation, nitrification, and denitrification) strongly control nitrogen dynamics in the soil. These reactions typically result in N enrichment of the substrate and depletion of the product. The heavy isotope N was first discovered by S. M. Naudé in 1929, soon after heavy isotopes of the neighbouring elements oxygen and carbon were discovered. It presents one of the lowest thermal neutron capture cross-sections of all isotopes. It is frequently used in nuclear magnetic resonance (NMR) spectroscopy to determine the structures of nitrogen-containing molecules, due to its fractional nuclear spin of one-half, which offers advantages for NMR such as narrower line width.
Nitrogen | Isotopes
All the houses were built . The exact dates and builders are not known.
Apache Nitrogen Products | Contributing properties
The difference in bond energy varies from 0.83 eV per atom in nitrogen gas to 4.94 eV per atom, so having a difference in energy of over 4 eV per atom. This cubic-gauche nitrogen is the highest energy non-nuclear material and is being investigated for use in explosives and rocket fuel. Its energy density is 33 kJ g which is over three times the energy density of HMX. cg-N has all bonds the same length of 1.346 Å at 115 GPa. The cubic-gauche structure for nitrogen is predicted to have bond lengths of 1.40 Å, bond angles of 114.0° and dihedral angles of −106.8°.
Solid nitrogen | Cubic gauche
In 1932, Chadwick exposed various elements, such as hydrogen and nitrogen, to the mysterious "beryllium radiation", and by measuring the energies of the recoiling charged particles, he deduced that the radiation was actually composed of electrically neutral particles which could not be massless like the gamma ray, but instead were required to have a mass similar to that of a proton. Chadwick now claimed these particles as Rutherford's neutrons. For his discovery of the neutron, Chadwick received the Nobel Prize in 1935.
Atomic theory | Discovery of nuclear particles
These results refined Ernest Rutherford's and Antonius van den Broek's model, which proposed that the atom contains in its nucleus a number of positive nuclear charges that is equal to its (atomic) number in the periodic table. Until these experiments, atomic number was not known to be a physical and experimental quantity. That it is equal to the atomic nuclear charge remains the accepted atomic model today. Chemical bonds between atoms were explained by Gilbert Newton Lewis in 1916, as the interactions between their constituent electrons. As the chemical properties of the elements were known to largely repeat themselves according to the periodic law, in 1919 the American chemist Irving Langmuir suggested that this could be explained if the electrons in an atom were connected or clustered in some manner.
Atom | Bohr model
by chemists of the time. The extension of this idea to substances in general necessarily led him to the law of multiple proportions, and the comparison with experiment brilliantly confirmed his deduction. In the paper "On the Proportion of the Several Gases in the Atmosphere", read by him in November 1802, the law of multiple proportions appears to be anticipated in the words: But there is reason to suspect that this sentence may have been added some time after the reading of the paper, which was not published until 1805.
John Dalton | Atomic weights
References to the word "neutron" in connection with the atom can be found in the literature as early as 1899, however.
Discovery of the neutron | Rutherford atom
He also coined the terms alpha, beta and gamma rays, he converted nitrogen into oxygen, and most importantly he supervised the students who did the Geiger–Marsden experiment (gold foil experiment) which showed that the 'plum pudding model' of the atom was wrong. In the plum pudding model, proposed by J. J. Thomson in 1904, the atom is composed of electrons surrounded by a 'cloud' of positive charge to balance the electrons' negative charge. To Rutherford, the gold foil experiment implied that the positive charge was confined to a very small nucleus leading first to the Rutherford model, and eventually to the Bohr model of the atom, where the positive nucleus is surrounded by the negative electrons.
Nuclear chemistry | History
Nitrogen has two stable isotopes: N and N. The first is much more common, making up 99.634% of natural nitrogen, and the second (which is slightly heavier) makes up the remaining 0.366%. This leads to an atomic weight of around 14.007 u. Both of these stable isotopes are produced in the CNO cycle in stars, but N is more common as its neutron capture is the rate-limiting step. N is one of the five stable odd–odd nuclides (a nuclide having an odd number of protons and neutrons); the other four are H, Li, B, and Ta. The relative abundance of N and N is practically constant in the atmosphere but can vary elsewhere, due to natural isotopic fractionation from biological redox reactions and the evaporation of natural ammonia or nitric acid.
Nitrogen | Isotopes
It shows 27% hydrogen at its surface. It is estimated to have had an initial mass of several million years ago.
WR 156 | Physical properties
Rutherford and others had noted the disparity between the mass of an atom, computed in atomic mass units, and the approximate charge required on the nucleus for the Rutherford model to work. The required charge of the atomic nucleus was usually about half its atomic mass. Antonius van den Broek boldly hypothesized that the required charge, denoted by "Z", was not half of the atomic weight for elements, but instead was exactly equal to the element's ordinal position in the periodic table. At that time, the positions of the elements in the periodic table were not known to have any physical significance.
Discovery of the neutron | Atomic number and Moseley's law
The number of neutrons in a nucleus usually has very little effect on an element's chemical properties (except in the case of hydrogen and deuterium). Thus, all carbon isotopes have nearly identical chemical properties because they all have six protons and six electrons, even though carbon atoms may, for example, have 6 or 8 neutrons. That is why the atomic number, rather than mass number or atomic weight, is considered the identifying characteristic of a chemical element. The symbol for atomic number is "Z".
Chemical element | Atomic number
Nitrogen compounds have a very long history, ammonium chloride having been known to Herodotus. They were well known by the Middle Ages. Alchemists knew nitric acid as "aqua fortis" (strong water), as well as other nitrogen compounds such as ammonium salts and nitrate salts. The mixture of nitric and hydrochloric acids was known as "aqua regia" (royal water), celebrated for its ability to dissolve gold, the king of metals. The discovery of nitrogen is attributed to the Scottish physician Daniel Rutherford in 1772, who called it "noxious air". Though he did not recognise it as an entirely different chemical substance, he clearly distinguished it from Joseph Black's "fixed air", or carbon dioxide.
Nitrogen | History
If the elements were ordered based on increasing atomic mass, however, periodicity in chemical properties was exhibited. Exceptions to this periodicity were apparent, however, such as cobalt and nickel. At the University of Manchester in 1913 Henry Moseley discussed the new Bohr model of the atom with the visiting Bohr. The model accounted for the electromagnetic emission spectrum from the hydrogen atom, and Moseley and Bohr wondered if the electromagnetic emission spectra of heavier elements such as cobalt and nickel would follow their ordering by weight, or by their position in the periodic table. In 1913-1914 Moseley tested the question experimentally by using X-ray diffraction techniques.
Discovery of the neutron | Atomic number and Moseley's law
The high compressive strength of 34 GPa, predicted for ideal crystals, is never obtained in the real bulk material due to the formation and movement of dislocations. However, it has been reached in Ni nanoparticles.
Nickel | Atomic and physical properties
The following table lists properties of selected radionuclides illustrating the range of properties and uses. Key: "Z" = atomic number; "N" = neutron number; DM = decay mode; DE = decay energy; EC = electron capture
Radionuclide | Examples
This is a list of chemical elements and their atomic properties, ordered by Atomic number. Since valence electrons are not clearly defined for the d-block and f-block elements, there not being a clear point at which further ionisation becomes unprofitable, a purely formal definition as number of electrons in the outermost shell has been used.
List of elements by atomic properties | Introduction
The nitrogen rule states that organic molecules that contain hydrogen, carbon, nitrogen, oxygen, silicon, phosphorus, sulfur, and the halogens have an odd nominal mass if they have an odd number of nitrogen atoms or an even mass if they have an even number of nitrogen atoms are present. The nitrogen rule is true for structures in which all of the atoms in the molecule have a number of covalent bonds equal to their standard valency, counting each sigma bond and pi bond as a separate covalent bond.
Mass spectral interpretation | Nitrogen rule
After graduating from Magdalene College in 1921, Blackett spent ten years working at the Cavendish Laboratory as an experimental physicist with Ernest Rutherford and in 1923 became a fellow of King's College, Cambridge, a position he held until 1933. Rutherford had found out that the nucleus of the nitrogen atom could be disintegrated by firing fast alpha particles into nitrogen. He asked Blackett to use a cloud chamber to find visible tracks of this disintegration, and by 1925, he had taken 23,000 photographs showing 415,000 tracks of ionized particles. Eight of these were forked, and this showed that the nitrogen atom-alpha particle combination had formed an atom of fluorine, which then disintegrated into an isotope of oxygen and a proton.
Patrick Blackett | Career and research
Table 3. General and Physical Properties of arthurite
Arthurite | Physical Properties
He realized that the plum pudding model could not be accurate and that the deflections of the alpha particles could only be explained if the positive and negative charges were separated from each other and that the mass of the atom was a concentrated point of positive charge. This justified the idea of a nuclear atom with a dense center of positive charge and mass.
Atomic nucleus | History
Under pressures higher than 110 GPa and temperatures around 2000 K nitrogen forms a network solid, bound by single covalent bonds in what is called a cubic-gauche structure, abbreviated as cg-N. This substance is very stiff with a bulk modulus around 298 GPa, similar to diamond. It is very high-energy. The cubic-gauche form has space group "I"23. The unit cell edge is 3.805 Å. There are eight nitrogen atoms per unit cell. The bond angles are very close to tetrahedral. The structure contains rings of nitrogen atoms that are fused together. The position of the lone pairs of electrons is ranged so that their overlap is minimised.
Solid nitrogen | Cubic gauche
Each element has a specific set of chemical properties as a consequence of the number of electrons present in the neutral atom, which is "Z" (the atomic number). The configuration of these electrons follows from the principles of quantum mechanics. The number of electrons in each element's electron shells, particularly the outermost valence shell, is the primary factor in determining its chemical bonding behavior. Hence, it is the atomic number alone that determines the chemical properties of an element; and it is for this reason that an element can be defined as consisting of "any" mixture of atoms with a given atomic number.
Atomic number | Chemical properties
Nitrogen bonds to almost all the elements in the periodic table except the first three noble gases, helium, neon, and argon, and some of the very short-lived elements after bismuth, creating an immense variety of binary compounds with varying properties and applications. Many binary compounds are known: with the exception of the nitrogen hydrides, oxides, and fluorides, these are typically called nitrides. Many stoichiometric phases are usually present for most elements (e.g. MnN, MnN, MnN, MnN, MnN, and MnN for 9.2 < "x" < 25.3). They may be classified as "salt-like" (mostly ionic), covalent, "diamond-like", and metallic (or interstitial), although this classification has limitations generally stemming from the continuity of bonding types instead of the discrete and separate types that it implies.
Nitrogen | Nitrides, azides, and nitrido complexes
With the development of quantum mechanics and experimental findings (such as the two slit diffraction of electrons), it was found that the orbiting electrons around a nucleus could not be fully described as particles, but needed to be explained by the wave-particle duality. In this sense, the electrons have the following properties: Wave-like properties: Particle-like properties:
Atomic orbital | Electron properties
Natural nitrogen (N) consists of two stable isotopes: the vast majority (99.6%) of naturally occurring nitrogen is nitrogen-14, with the remainder being nitrogen-15. Fourteen radioisotopes are also known, with atomic masses ranging from 10 to 25, along with one nuclear isomer, N. All of these radioisotopes are short-lived, the longest-lived being nitrogen-13 with a half-life of 9.965 minutes. All of the others have half-lives below 7.15 seconds, with most of these being below 620 milliseconds. Most of the isotopes with atomic mass numbers below 14 decay to isotopes of carbon, while most of the isotopes with masses above 15 decay to isotopes of oxygen.
Isotopes of nitrogen | Introduction
The term "gauche" refers to the odd dihedral angles, if it were 0° it would be called "cis", and if 180° it would be called "trans". The dihedral angle Φ is related to the bond angle θ by sec(Φ) = sec(θ) − 1. The coordinate of one atom in the unit cell at x,x,x also determines the bond angle by cos(θ) = x(x-1/4)/(x+(x-1/4)).
Solid nitrogen | Cubic gauche
A helium nucleus was presumed to be composed of four protons plus two "nuclear electrons" (electrons bound inside the nucleus) to cancel two of the charges. At the other end of the periodic table, a nucleus of gold with a mass 197 times that of hydrogen was thought to contain 118 nuclear electrons in the nucleus to give it a residual charge of +79, consistent with its atomic number.
Atomic number | The proton and the idea of nuclear electrons
This adduct was first reported by Bernard Courtois in 1812, and its formula was finally determined in 1905 by Oswald Silberrad. Its solid state structure consists of chains of -NI-I-NI-I-NI-I-... Ammonia molecules are situated between the chains. When kept cold in the dark and damp with ammonia, NI · NH is stable.
Nitrogen triiodide | Structure of NI and its derivatives
During an analysis of the uranium containing mineral uranite he discovered that a gas evolved. He identified this gas by spectroscopic methods to be nitrogen. Several years later in 1895 William Ramsay did similar experiments with uranium containing minerals and discovered by similar methods that the gas was a mixture of argon and helium which until then had only been detected in the corona of stars. A reexamination by Ramsay of Hillebrand's samples showed that the gas from uranite contained a large amount of nitrogen.
William Francis Hillebrand | Studies of uranium
Hydrogen Helium Carbon
Properties of metals, metalloids and nonmetals | Nonmetals
With dideuterium, a clathrate (N)D exits around 70 GPa.
Solid nitrogen | Hydrogen
The nitrogen-vacancy center is a point defect in the diamond lattice. It consists of a nearest-neighbor pair of a nitrogen atom, which substitutes for a carbon atom, and a lattice vacancy. Two charge states of this defect, neutral N-V and negative N-V, are known from spectroscopic studies using optical absorption, photoluminescence (PL), electron paramagnetic resonance (EPR) and optically detected magnetic resonance (ODMR), which can be viewed as a hybrid of PL and EPR; most details of the structure originate from EPR. A nitrogen atom has five valence electrons. Three of them covalently bond to the carbon atoms and two remain non-bonded and are called a lone pair.
Nitrogen-vacancy center | Structure
The chemical properties of rutherfordium are characterized only partly. They compare well with the chemistry of the other group 4 elements, even though some calculations had indicated that the element might show significantly different properties due to relativistic effects. In the 1960s, small amounts of rutherfordium were produced in the Joint Institute for Nuclear Research in the Soviet Union and at Lawrence Berkeley National Laboratory in California. The priority of the discovery and therefore the naming of the element was disputed between Soviet and American scientists, and it was not until 1997 that International Union of Pure and Applied Chemistry (IUPAC) established rutherfordium as the official name for the element.
Rutherfordium | Introduction
As a result, he discovered the emission of a subatomic particle which, in 1919, he called the "hydrogen atom" but, in 1920, he more accurately named the proton.
Ernest Rutherford | Introduction
The fact that there was a component of air that does not support combustion was clear to Rutherford, although he was not aware that it was an element. Nitrogen was also studied at about the same time by Carl Wilhelm Scheele, Henry Cavendish, and Joseph Priestley, who referred to it as "burnt air" or "phlogisticated air". Nitrogen gas was inert enough that Antoine Lavoisier referred to it as "mephitic air" or "azote", from the Greek word (azotikos), "no life". In an atmosphere of pure nitrogen, animals died and flames were extinguished. Though Lavoisier's name was not accepted in English, since it was pointed out that almost all gases (indeed, with the sole exception of oxygen) are mephitic, it is used in many languages (French, Italian, Portuguese, Polish, Russian, Albanian, Turkish, etc.; the German "Stickstoff" similarly refers to the same characteristic, viz.
Nitrogen | History
In 1919, Ernest Rutherford was able to accomplish transmutation of nitrogen into oxygen at the University of Manchester, using alpha particles directed at nitrogen N + α → O + p. This was the first observation of an induced nuclear reaction, that is, a reaction in which particles from one decay are used to transform another atomic nucleus. Eventually, in 1932 at Cambridge University, a fully artificial nuclear reaction and nuclear transmutation was achieved by Rutherford's colleagues John Cockcroft and Ernest Walton, who used artificially accelerated protons against lithium-7, to split the nucleus into two alpha particles. The feat was popularly known as "splitting the atom", although it was not the modern nuclear fission reaction later discovered in heavy elements, in 1938 by the German scientists Otto Hahn, Lise Meitner, and Fritz Strassmann.
Nuclear reaction | History
The atomic nucleus is the small, dense region consisting of protons and neutrons at the center of an atom, discovered in 1911 by Ernest Rutherford based on the 1909 Geiger–Marsden gold foil experiment. After the discovery of the neutron in 1932, models for a nucleus composed of protons and neutrons were quickly developed by Dmitri Ivanenko and Werner Heisenberg. An atom is composed of a positively-charged nucleus, with a cloud of negatively-charged electrons surrounding it, bound together by electrostatic force. Almost all of the mass of an atom is located in the nucleus, with a very small contribution from the electron cloud.
Atomic nucleus | Introduction
Hydrogen: Cavendish, in 1766, was the first to distinguish hydrogen from other gases, although Paracelsus around 1500, Robert Boyle (1670), and Joseph Priestley (?) had observed its production by reacting strong acids with metals. Lavoisier named it in 1793. Oxygen: Carl Wilhelm Scheele obtained oxygen by heating mercuric oxide and nitrates in 1771, but did not publish his findings until 1777. Priestley also prepared this new "air" by 1774, but only Lavoisier recognized it as a true element; he named it in 1777. Nitrogen: Rutherford discovered nitrogen while he was studying at the University of Edinburgh. He showed that the air in which animals breathed, after removal of exhaled carbon dioxide, was no longer able to burn a candle.
Nonmetal | 18th century: H, O, N, (Te), Cl
Prout's hypothesis was an influence on Ernest Rutherford when he succeeded in "knocking" hydrogen nuclei out of nitrogen atoms with alpha particles in 1917, and thus concluded that perhaps the nuclei of all elements were made of such particles (the hydrogen nucleus), which in 1920 he suggested be named protons, from the suffix for particles, added to the stem of Prout's word "protyle". The assumption as discussed by Rutherford was of a nucleus consisting of Z+N=A protons plus N electrons somehow trapped within thereby reducing the positive charge to +Z as observed and vaguely explaining beta decay radioactivity. Such a nuclear constitution was known to be inconsistent with dynamics either classical or early quantum but seemed inevitable until the neutron hypothesis by Rutherford and discovery by Chadwick.
Prout's hypothesis | Introduction
The discovery of the neutron and its properties was central to the extraordinary developments in atomic physics in the first half of the 20th century. Early in the century, Ernest Rutherford developed a crude model of the atom, based on the gold foil experiment of Hans Geiger and Ernest Marsden. In this model, atoms had their mass and positive electric charge concentrated in a very small nucleus. By 1920 chemical isotopes had been discovered, the atomic masses had been determined to be (approximately) integer multiples of the mass of the hydrogen atom, and the atomic number had been identified as the charge on the nucleus.
Discovery of the neutron | Introduction
Rutherford remarked that: "Hahn has a special nose for discovering new elements."
Otto Hahn | Discovery of radio thorium, and other "new elements"
This stands in contrast to the well-characterised cationic and anionic polynitrogens azide (), pentazenium (), and pentazolide (cyclic aromatic ). Under extremely high pressures (1.1 million atm) and high temperatures (2000 K), as produced in a diamond anvil cell, nitrogen polymerises into the single-bonded cubic gauche crystal structure. This structure is similar to that of diamond, and both have extremely strong covalent bonds, resulting in its nickname "nitrogen diamond". At atmospheric pressure, molecular nitrogen condenses (liquefies) at 77 K (−195.79 °C) and freezes at 63 K (−210.01 °C) into the beta hexagonal close-packed crystal allotropic form. Below 35.4 K (−237.6 °C) nitrogen assumes the cubic crystal allotropic form (called the alpha phase).
Nitrogen | Allotropes
One of the first studies of condensed states of matter was by English chemist Humphry Davy, in the first decades of the nineteenth century. Davy observed that of the forty chemical elements known at the time, twenty-six had metallic properties such as lustre, ductility and high electrical and thermal conductivity. This indicated that the atoms in John Dalton's atomic theory were not indivisible as Dalton claimed, but had inner structure. Davy further claimed that elements that were then believed to be gases, such as nitrogen and hydrogen could be liquefied under the right conditions and would then behave as metals.
Condensed matter physics | Classical physics
The neutron is a subatomic particle, symbol or , which has a neutral (not positive or negative) charge and a mass slightly greater than that of a proton. Protons and neutrons constitute the nuclei of atoms. Since protons and neutrons behave similarly within the nucleus, and each has a mass of approximately one atomic mass unit, they are both referred to as nucleons. Their properties and interactions are described by nuclear physics. The chemical properties of an atom are mostly determined by the configuration of electrons that orbit the atom's heavy nucleus. The electron configuration is determined by the charge of the nucleus, set by the number of protons, or atomic number.
Neutron | Introduction
No chemical analysis was possible at the time, but helium was later found to be a noble gas. Before them, in 1784, the English chemist and physicist Henry Cavendish had discovered that air contains a small proportion of a substance less reactive than nitrogen. A century later, in 1895, Lord Rayleigh discovered that samples of nitrogen from the air were of a different density than nitrogen resulting from chemical reactions. Along with Scottish scientist William Ramsay at University College, London, Lord Rayleigh theorized that the nitrogen extracted from air was mixed with another gas, leading to an experiment that successfully isolated a new element, argon, from the Greek word (, "idle" or "lazy").
Noble gas | History
By definition, any two atoms with an identical number of "protons" in their nuclei belong to the same chemical element. Atoms with equal numbers of protons but a different number of "neutrons" are different isotopes of the same element. For example, all hydrogen atoms admit exactly one proton, but isotopes exist with no neutrons (hydrogen-1, by far the most common form, also called protium), one neutron (deuterium), two neutrons (tritium) and more than two neutrons. The known elements form a set of atomic numbers, from the single-proton element hydrogen up to the 118-proton element oganesson. All known isotopes of elements with atomic numbers greater than 82 are radioactive, although the radioactivity of element 83 (bismuth) is so slight as to be practically negligible.
Atom | Nuclear properties
The table below is a summary of the key physical and atomic properties of the halogens. Data marked with question marks are either uncertain or are estimations partially based on periodic trends rather than observations.
Halogen | Physical and atomic
Such a demonstration must establish properties, either physical or chemical, of the new element and establish that they are those of a previously unknown element. The main techniques used to demonstrate atomic number are cross-reactions (creating claimed nuclides as parents or daughters of other nuclides produced by a different reaction) and anchoring decay chains to known daughter nuclides. For the JWP, priority in confirmation takes precedence over the date of the original claim. Both teams set out to confirm their results by these methods.
Nihonium | Road to confirmation
J. Thomson. Wilhelm Wien in 1898 identified the hydrogen ion as the particle with the highest charge-to-mass ratio in ionized gases. Following the discovery of the atomic nucleus by Ernest Rutherford in 1911, Antonius van den Broek proposed that the place of each element in the periodic table (its atomic number) is equal to its nuclear charge. This was confirmed experimentally by Henry Moseley in 1913 using X-ray spectra. In 1917 (in experiments reported in 1919 and 1925), Rutherford proved that the hydrogen nucleus is present in other nuclei, a result usually described as the discovery of protons. These experiments began after Rutherford had noticed that, when alpha particles were shot into air (mostly nitrogen), his scintillation detectors showed the signatures of typical hydrogen nuclei as a product.
Proton | History
Solid nitrogen is the solid form of the element nitrogen. It is an important component of the surfaces of Pluto and outer moons of the Solar System such as Neptune's Triton. Under low or moderate pressure solid nitrogen contains dinitrogen molecules held together by London dispersion forces. Non-molecular forms of solid nitrogen produced by extreme pressures have a higher energy density than any other non-nuclear material. Solid nitrogen was first made in 1884, by first liquefying hydrogen with evaporating liquid nitrogen, and then allowing the liquid hydrogen to freeze the nitrogen. Karol Olszewski achieved a world record lowest temperature in 1884 by evaporating vapour from solid nitrogen getting down to 48 K.
Solid nitrogen | Introduction
This would make the nuclear charge nearly equal to atomic number in smaller atoms, with some deviation from this rule for the largest atoms, such as gold. For example, Rutherford found the charge on gold to be about 100 units and thought perhaps that it might be exactly 98 (which would be close to half its atomic weight). But gold's place in the periodic table (and thus its atomic number) was known to be 79. Thus Rutherford did not make the proposal that the number of charges in the nucleus of an atom might be "exactly" equal to its place on the periodic table (atomic number).
Antonius van den Broek | Major contribution to science
Zero has been proposed as the atomic number of the theoretical element tetraneutron. It has been shown that a cluster of four neutrons may be stable enough to be considered an atom in its own right. This would create an element with no protons and no charge on its nucleus. As early as 1926, Andreas von Antropoff coined the term neutronium for a conjectured form of matter made up of neutrons with no protons, which he placed as the chemical element of atomic number zero at the head of his new version of the periodic table. It was subsequently placed as a noble gas in the middle of several spiral representations of the periodic system for classifying the chemical elements.
0 | Chemistry
In 1915, the reason for nuclear charge being quantized in units of "Z", which were now recognized to be the same as the element number, was not understood. An old idea called Prout's hypothesis had postulated that the elements were all made of residues (or "protyles") of the lightest element hydrogen, which in the Bohr-Rutherford model had a single electron and a nuclear charge of one. However, as early as 1907, Rutherford and Thomas Royds had shown that alpha particles, which had a charge of +2, were the nuclei of helium atoms, which had a mass four times that of hydrogen, not two times.
Atomic number | The proton and the idea of nuclear electrons
All consideration of nuclear electrons ended with James Chadwick's discovery of the neutron in 1932. An atom of gold now was seen as containing 118 neutrons rather than 118 nuclear electrons, and its positive charge now was realized to come entirely from a content of 79 protons. After 1932, therefore, an element's atomic number "Z" was also realized to be identical to the proton number of its nuclei.
Atomic number | The discovery of the neutron makes "Z" the proton number
Solid nitrogen has several properties relevant to its formation of rocks in the outer Solar System. Even at the low temperatures of solid nitrogen it is fairly volatile and can sublime to form an atmosphere, or condense back into nitrogen frost. At 58 K the ultimate compressive strength is 0.24 MPa. Strength increases as temperature lowers becoming 0.54 MPa at 40.6 K. Elastic modulus varies from 161 to 225 MPa over the same range. Compared to other materials, solid nitrogen loses cohesion at low pressures and flows in the form of glaciers when amassed. Yet its density is higher than that of water ice, so the forces of buoyancy will naturally transport blocks of water ice towards the surface.
Solid nitrogen | Bulk properties
If an atom has more or fewer electrons than its atomic number, then it becomes respectively negatively or positively charged as a whole; a charged atom is called an ion. Electrons have been known since the late 19th century, mostly thanks to J.J. Thomson; see history of subatomic physics for details. Protons have a positive charge and a mass 1,836 times that of the electron, at . The number of protons in an atom is called its atomic number. Ernest Rutherford (1919) observed that nitrogen under alpha-particle bombardment ejects what appeared to be hydrogen nuclei. By 1920 he had accepted that the hydrogen nucleus is a distinct particle within the atom and named it proton.
Atom | Subatomic particles
The isotope of nitrogen had an extra neutron in the nucleus, which made it heavier.
Meselson–Stahl experiment | Introduction