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The shapes of the five orbitals occupied in nitrogen. The two colours show the phase or sign of the wave function in each region. From left to right: 1s, 2s (cutaway to show internal structure), 2p x, 2p y, 2p z. A nitrogen atom has seven electrons. In the ground state, they are arranged in the electron configuration 1s 2 2s 2 2p 1 x2p 1 y2p 1 z. It therefore has five in the 2s and 2p orbitals, three of which (the p-electrons) are unpaired. It has one of the highest among the elements (3.04 on the Pauling scale), exceeded only by (3.16), (3.44), and (3.98).
Following periodic trends, its single-bond of 71 pm is smaller than those of (84 pm) and (76 pm), while it is larger than those of oxygen (66 pm) and fluorine (57 pm). The anion, N 3−, is much larger at 146 pm, similar to that of the (O 2−: 140 pm) and (F −: 133 pm) anions. The first three ionisation energies of nitrogen are 1.402, 2.856, and 4.577 MJmol −1, and the sum of the fourth and fifth is 16.920 MJmol −1. 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, especially in nitrogen, oxygen, and fluorine. The 2p subshell is very small and has a very similar radius to the 2s shell, facilitating. 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.
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 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, and. 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 nitrogen shows some similarities to the period 3 element in the next group from magnesium to sulfur (known as the ), their degree drops off quite abruptly past the boron–silicon pair, so that the similarities of nitrogen to sulfur are mostly limited to sulfur nitride ring compounds when both elements are the only ones present. Nitrogen resembles oxygen far more than it does carbon with its high electronegativity and concomitant capability for and the ability to form by donating its of electrons.
It does not share carbon's proclivity for, with the longest chain of nitrogen yet discovered being composed of only eight nitrogen atoms (PhN=N–N(Ph)–N=N–N(Ph)–N=NPh). One property nitrogen does share with both its horizontal neighbours is its preferentially forming multiple bonds, typically with carbon, nitrogen, or oxygen atoms, through p π–p π interactions; thus, for example, nitrogen occurs as diatomic molecules and thus has very much lower (−210 °C) and (−196 °C) than the rest of its group, as the N 2 molecules are only held together by weak and there are very few electrons available to create significant instantaneous dipoles. This is not possible for its vertical neighbours; thus, the, -, -, and -compounds, and -derivatives find no echo with phosphorus, arsenic, antimony, or bismuth. By the same token, however, the complexity of the phosphorus oxoacids finds no echo with nitrogen. Of dinitrogen molecule, N 2. There are five bonding orbitals and two antibonding orbitals (marked with an asterisk), giving a total bond order of three.
Atomic nitrogen, also known as active nitrogen, is highly reactive, being a 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 2 molecule, they may release so much energy on collision with even such stable molecules as and 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. Given the great reactivity of atomic nitrogen, elemental nitrogen usually occurs as molecular N 2,. This molecule is a colourless, odourless, and tasteless gas at standard conditions: it melts at −210 °C and boils at −196 °C. Dinitrogen is mostly unreactive at room temperature, but it will nevertheless react with metal and some complexes. This is due to its bonding, which is unique among the diatomic elements at standard conditions in that it has an N≡N.
Triple bonds have short bond lengths (in this case, 109.76 pm) and high dissociation energies (in this case, 945.41 kJ/mol), and are thus very strong, explaining dinitrogen's chemical inertness. There are some theoretical indications that other nitrogen oligomers and polymers may be possible. If they could be synthesised, they may have potential applications as materials with a very high energy density, that could be used as powerful propellants or explosives. This is because they should all decompose to dinitrogen, whose N≡N triple bond (bond energy 946 kJ⋅mol −1) is much stronger than those of the N=N double bond (418 kJ⋅mol −1) or the N–N single bond (160 kJ⋅mol −1): indeed, the triple bond has more than thrice the energy of the single bond.
(The opposite is true for the heavier pnictogens, which prefer polyatomic allotropes.) A great disadvantage is that most neutral polynitrogens are not expected to have a large barrier towards decomposition, and that the few exceptions would be even more challenging to synthesise than the long-sought but still unknown. This stands in contrast to the well-characterised cationic and anionic polynitrogens ( N − 3), ( N + 5), and (cyclic aromatic N − 5). Under extremely high pressures (1.1 million ) and high temperatures (2000 K), as produced in a, nitrogen polymerises into the single-bonded cubic gauche crystal structure. This structure is similar to that of, and both have extremely strong, resulting in its nickname 'nitrogen diamond'.
At, molecular nitrogen at 77 (−195.79 °) and at 63 K (−210.01 °C) into the beta crystal form. Below 35.4 K (−237.6 °C) nitrogen assumes the crystal allotropic form (called the alpha phase)., a colourless fluid resembling water in appearance, but with 80.8% of the density (the density of liquid nitrogen at its boiling point is 0.808 g/mL), is a common. Has many crystalline modifications. It forms a significant dynamic surface coverage on and outer moons of the Solar System such as.
Even at the low temperatures of solid nitrogen it is fairly volatile and can to form an atmosphere, or condense back into nitrogen frost. It is very weak and flows in the form of glaciers and on Triton of nitrogen gas come from the polar ice cap region. Dinitrogen complexes. Standard reduction potentials for nitrogen-containing species. Top diagram shows potentials at pH 0; bottom diagram shows potentials at pH 14. Industrially, (NH 3) is the most important compound of nitrogen and is prepared in larger amounts than any other compound, because it contributes significantly to the nutritional needs of terrestrial organisms by serving as a precursor to food and fertilisers.
It is a colourless alkaline gas with a characteristic pungent smell. The presence of has very significant effects on ammonia, conferring on it its high melting (−78 °C) and boiling (−33 °C) points. As a liquid, it is a very good solvent with a high heat of vaporisation (enabling it to be used in vacuum flasks), that also has a low viscosity and electrical conductivity and high, and is less dense than water. However, the hydrogen bonding in NH 3 is weaker than that in H 2O due to the lower electronegativity of nitrogen compared to oxygen and the presence of only one lone pair in NH 3 rather than two in H 2O. It is a weak base in aqueous solution ( 4.74); its conjugate acid is, NH + 4. It can also act as an extremely weak acid, losing a proton to produce the amide anion, NH − 2.
It thus undergoes self-dissociation, similar to water, to produce ammonium and amide. Ammonia burns in air or oxygen, though not readily, to produce nitrogen gas; it burns in fluorine with a greenish-yellow flame to give. Reactions with the other nonmetals are very complex and tend to lead to a mixture of products.
Ammonia reacts on heating with metals to give nitrides. Many other binary nitrogen hydrides are known, but the most important are (N 2H 4) and (HN 3). Although it is not a nitrogen hydride, (NH 2OH) is similar in properties and structure to ammonia and hydrazine as well. Hydrazine is a fuming, colourless liquid that smells similarly to ammonia. Its physical properties are very similar to those of water (melting point 2.0 °C, boiling point 113.5 °C, density 1.00 gcm −3). Despite it being an endothermic compound, it is kinetically stable. It burns quickly and completely in air very exothermically to give nitrogen and water vapour.
It is a very useful and versatile reducing agent and is a weaker base than ammonia. It is also commonly used as a rocket fuel. Hydrazine is generally made by reaction of ammonia with alkaline in the presence of gelatin or glue: NH 3 + OCl − → NH 2Cl + OH − NH 2Cl + NH 3 → N 2H + 5 + Cl − (slow) N 2H + 5 + OH − → N 2H 4 + H 2O (fast) (The attacks by hydroxide and ammonia may be reversed, thus passing through the intermediate NHCl − instead.) The reason for adding gelatin is that it removes metal ions such as Cu 2+ that catalyses the destruction of hydrazine by reaction with (NH 2Cl) to produce and nitrogen. (HN 3) was first produced in 1890 by the oxidation of aqueous hydrazine by nitrous acid.
It is very explosive and even dilute solutions can be dangerous. It has a disagreeable and irritating smell and is a potentially lethal (but not cumulative) poison. It may be considered the conjugate acid of the azide anion, and is similarly analogous to the. Halides and oxohalides.
All four simple nitrogen trihalides are known. A few mixed halides and hydrohalides are known, but are mostly unstable and uninteresting: examples include NClF 2, NCl 2F, NBrF 2, NF 2H, NCl 2H, and NClH 2. Five nitrogen fluorides are known. (NF 3, first prepared in 1928) is a colourless and odourless gas that is thermodynamically stable, and most readily produced by the electrolysis of molten dissolved in anhydrous. Like, it is not at all reactive and is stable in water or dilute aqueous acids or alkalis. Only when heated does it act as a fluorinating agent, and it reacts with, arsenic, antimony, and bismuth on contact at high temperatures to give (N 2F 4).
The cations NF + 4 and N 2F + 3 are also known (the latter from reacting tetrafluorohydrazine with strong fluoride-acceptors such as ), as is ONF 3, which has aroused interest due to the short N–O distance implying partial double bonding and the highly polar and long N–F bond. Tetrafluorohydrazine, unlike hydrazine itself, can dissociate at room temperature and above to give the radical NF 2. (FN 3) is very explosive and thermally unstable. (N 2F 2) exists as thermally interconvertible cis and trans isomers, and was first found as a product of the thermal decomposition of FN 3.
(NCl 3) is a dense, volatile, and explosive liquid whose physical properties are similar to those of, although one difference is that NCl 3 is easily hydrolysed by water while CCl 4 is not. It was first synthesised in 1811 by, who lost three fingers and an eye to its explosive tendencies. As a dilute gas it is less dangerous and is thus used industrially to bleach and sterilise flour. (NBr 3), first prepared in 1975, is a deep red, temperature-sensitive, volatile solid that is explosive even at −100 °C.
(NI 3) is still more unstable and was only prepared in 1990. Its adduct with ammonia, which was known earlier, is very shock-sensitive: it can be set off by the touch of a feather, shifting air currents, or even. For this reason, small amounts of nitrogen triiodide are sometimes synthesised as a demonstration to high school chemistry students or as an act of 'chemical magic'. Two series of nitrogen oxohalides are known: the nitrosyl halides (XNO) and the nitryl halides (XNO 2). The first are very reactive gases that can be made by directly halogenating nitrous oxide. (NOF) is colourless and a vigorous fluorinating agent. (NOCl) behaves in much the same way and has often been used as an ionising solvent.
(NOBr) is red. The reactions of the nitryl halides are mostly similar: (FNO 2) and (ClNO 2) are likewise reactive gases and vigorous halogenating agents. Nitrogen dioxide at −196 °C, 0 °C, 23 °C, 35 °C, and 50 °C. NO 2 converts to colourless dinitrogen tetroxide ( N 2O 4) at low temperatures, and reverts to NO 2 at higher temperatures. Nitrogen forms nine molecular oxides, some of which were the first gases to be identified: N 2O , NO , N 2O 3 , NO 2 , N 2O 4 , N 2O 5 , N 4O , and N(NO 2) 3.
All are thermally unstable towards decomposition to their elements. One other possible oxide that has not yet been synthesised is (N 4O), an aromatic ring. Nitrous oxide (N 2O), better known as laughing gas, is made by thermal decomposition of molten at 250 °C. This is a redox reaction and thus nitric oxide and nitrogen are also produced as byproducts.
It is mostly used as a propellant and aerating agent for, and was formerly commonly used as an anaesthetic. Despite appearances, it cannot be considered to be the of (H 2N 2O 2) because that acid is not produced by the dissolution of nitrous oxide in water. It is rather unreactive (not reacting with the halogens, the alkali metals, or at room temperature, although reactivity increases upon heating) and has the unsymmetrical structure N–N–O (N≡N +O −↔ −N=N +=O): above 600 °C it dissociates by breaking the weaker N–O bond.
Nitric oxide (NO) is the simplest stable molecule with an odd number of electrons. In mammals, including humans, it is an important cellular involved in many physiological and pathological processes. It is formed by catalytic oxidation of ammonia.
It is a colourless paramagnetic gas that, being thermodynamically unstable, decomposes to nitrogen and oxygen gas at 1100–1200 °C. Its bonding is similar to that in nitrogen, but one extra electron is added to a π.
antibonding orbital and thus the bond order has been reduced to approximately 2.5; hence dimerisation to O=N–N=O is unfavourable except below the boiling point (where the cis isomer is more stable) because it does not actually increase the total bond order and because the unpaired electron is delocalised across the NO molecule, granting it stability. There is also evidence for the asymmetric red dimer O=N–O=N when nitric oxide is condensed with polar molecules. It reacts with oxygen to give brown nitrogen dioxide and with halogens to give nitrosyl halides. It also reacts with transition metal compounds to give nitrosyl complexes, most of which are deeply coloured. Blue dinitrogen trioxide (N 2O 3) is only available as a solid because it rapidly dissociates above its melting point to give nitric oxide, nitrogen dioxide (NO 2), and dinitrogen tetroxide (N 2O 4). The latter two compounds are somewhat difficult to study individually because of the equilibrium between them. Although sometimes dinitrogen tetroxide can react by heterolytic fission to and in a medium with high dielectric constant.
Nitrogen dioxide is an acrid, corrosive brown gas. Both compounds may be easily prepared by decomposing a dry metal nitrate. Both react with water to form. Dinitrogen tetroxide is very useful for the preparation of anhydrous metal nitrates and nitrato complexes, and it became the storable oxidiser of choice for many rockets in both the United States and by the late 1950s.
This is because it is a in combination with a -based and can be easily stored since it is liquid at room temperature. The thermally unstable and very reactive dinitrogen pentoxide (N 2O 5) is the anhydride of, and can be made from it by dehydration with.
It is of interest for the preparation of explosives. It is a, colourless crystalline solid that is sensitive to light. In the solid state it is ionic with structure NO 2 +NO 3 −; as a gas and in solution it is molecular O 2N–O–NO 2.
Hydration to nitric acid comes readily, as does analogous reaction with giving (HOONO 2). It is a violent oxidising agent.
Gaseous dinitrogen pentoxide decomposes as follows: N 2O 5 ⇌ NO 2 + NO 3 → NO 2 + O 2 + NO N 2O 5 + NO ⇌ 3 NO 2 Oxoacids, oxoanions, and oxoacid salts Many nitrogen are known, though most of them are unstable as pure compounds and are known only as aqueous solution or as salts. (H 2N 2O 2) is a weak diprotic acid with the structure HON=NOH (p K a1 6.9, p K a2 11.6).
Acidic solutions are quite stable but above pH 4 base-catalysed decomposition occurs via HONNO − to nitrous oxide and the hydroxide anion. (involving the N 2O 2− 2 anion) are stable to reducing agents and more commonly act as reducing agents themselves. They are an intermediate step in the oxidation of ammonia to nitrite, which occurs in the. Hyponitrite can act as a bridging or chelating bidentate ligand. (HNO 2) is not known as a pure compound, but is a common component in gaseous equilibria and is an important aqueous reagent: its aqueous solutions may be made from acidifying cool aqueous ( NO − 2, bent) solutions, although already at room temperature disproportionation to and nitric oxide is significant. It is a weak acid with p K a 3.35 at 18 °C.
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They may be analysed by their oxidation to nitrate. They are readily reduced to nitrous oxide and nitric oxide by, to hyponitrous acid with (II), and to ammonia with.
Salts of N 2H + 5 react with nitrous acid to produce azides which further react to give nitrous oxide and nitrogen. Is mildly toxic in concentrations above 100 mg/kg, but small amounts are often used to cure meat and as a preservative to avoid bacterial spoilage. It is also used to synthesise hydroxylamine and to diazotise primary aromatic amines as follows: ArNH 2 + HNO 2 → ArNNCl + 2 H 2O Nitrite is also a common ligand that can coordinate in five ways. The most common are nitro (bonded from the nitrogen) and nitrito (bonded from an oxygen). Nitro-nitrito isomerism is common, where the nitrito form is usually less stable.
See also: Nitrogen is the most common pure element in the earth, making up 78.1% of the entire volume of the atmosphere. Despite this, it is not very abundant in Earth's crust, making up only 19 of this, on par with, and. The only important nitrogen minerals are (, saltpetre) and (, Chilean saltpetre). However, these have not been an important source of nitrates since the 1920s, when the industrial synthesis of ammonia and nitric acid became common. Nitrogen compounds constantly interchange between the atmosphere and living organisms. Nitrogen must first be processed, or ', into a plant-usable form, usually ammonia.
Some nitrogen fixation is done by lightning strikes producing the nitrogen oxides, but most is done by bacteria through enzymes known as (although today industrial nitrogen fixation to ammonia is also significant). Horoscope explorer software free download full version in bengali. When the ammonia is taken up by plants, it is used to synthesise proteins. These plants are then digested by animals who use the nitrogen compounds to synthesise their own proteins and excrete nitrogen–bearing waste. Finally, these organisms die and decompose, undergoing bacterial and environmental oxidation and, returning free dinitrogen to the atmosphere.
Industrial nitrogen fixation by the is mostly used as fertiliser, although excess nitrogen–bearing waste, when leached, leads to of freshwater and the creation of marine, as nitrogen-driven bacterial growth depletes water oxygen to the point that all higher organisms die. Furthermore, nitrous oxide, which is produced during denitrification, attacks the atmospheric. Many saltwater fish manufacture large amounts of to protect them from the high effects of their environment; conversion of this compound to is responsible for the early odour in unfresh saltwater fish. In animals, (derived from an ), serves as an important regulatory molecule for circulation. Nitric oxide's rapid reaction with water in animals results in production of its metabolite. Animal of nitrogen in proteins, in general, results in of, while animal metabolism of results in excretion of and. The characteristic odour of animal flesh decay is caused by the creation of long-chain, nitrogen-containing, such as and, which are breakdown products of the amino acids and, respectively, in decaying proteins.
Production Nitrogen gas is an produced by the fractional of liquid, or by mechanical means using gaseous air (pressurised reverse or ). Nitrogen gas generators using membranes or pressure swing adsorption (PSA) are typically more cost and energy efficient than bulk delivered nitrogen. Commercial nitrogen is often a byproduct of air-processing for industrial concentration of for steelmaking and other purposes. When supplied compressed in cylinders it is often called OFN (oxygen-free nitrogen). Commercial-grade nitrogen already contains at most 20 ppm oxygen, and specially purified grades containing at most 2 ppm oxygen and 10 ppm are also available. In a chemical laboratory, it is prepared by treating an aqueous solution of with.
NH 4Cl + NaNO 2 → N 2 + NaCl + 2 H 2O Small amounts of the impurities NO and HNO 3 are also formed in this reaction. The impurities can be removed by passing the gas through aqueous sulfuric acid containing. Very pure nitrogen can be prepared by the thermal decomposition of. 2 NaN 3 → 2 Na + 3 N 2 Applications Gas The applications of nitrogen compounds are naturally extremely widely varied due to the huge size of this class: hence, only applications of pure nitrogen itself will be considered here. Two-thirds of nitrogen produced by industry is sold as the gas and the remaining one-third as the liquid.
The gas is mostly used as an inert atmosphere whenever the oxygen in the air would pose a fire, explosion, or oxidising hazard. Some examples include:. As a, pure or mixed with, to nitrogenate and preserve the freshness of packaged or bulk foods (by delaying and other forms of ). Pure nitrogen as food additive is labeled in the with the E941. In as an inexpensive alternative to. In the manufacture of.
In the of steel. In some aircraft fuel systems to reduce fire hazard (see ). To inflate race car and aircraft, reducing the problems caused by moisture and in natural air. Nitrogen is commonly used during sample preparation in. It is used to concentrate and reduce the volume of liquid samples. Directing a pressurised stream of nitrogen gas perpendicular to the surface of the liquid causes the solvent to evaporate while leaving the solute(s) and un-evaporated solvent behind.
Nitrogen can be used as a replacement, or in combination with, to pressurise kegs of some, particularly and British, due to the smaller it produces, which makes the dispensed beer smoother and. A pressure-sensitive nitrogen capsule known commonly as a ' allows nitrogen-charged beers to be packaged in and. Nitrogen tanks are also replacing carbon dioxide as the main power source for. Nitrogen must be kept at higher pressure than CO 2, making N 2 tanks heavier and more expensive. Nitrogen gas has become the inert gas of choice for, and is under consideration as a replacement for lethal injection in. Nitrogen gas, formed from the decomposition of, is used for the inflation of. Air balloon submerged in liquid nitrogen Liquid nitrogen is a.
When insulated in proper containers such as, it can be transported without much. Like, the main use of liquid nitrogen is as a. Among other things, it is used in the of blood, reproductive cells ( and ), and other biological samples and materials.
It is used in the clinical setting in to remove cysts and warts on the skin. It is used in for certain laboratory equipment and to cool.
It has also been used to cool and other devices in computers that are, and that produce more heat than during normal operation. Other uses include freeze-grinding and machining materials that are soft or rubbery at room temperature, shrink-fitting and assembling engineering components, and more generally to attain very low temperatures whenever necessary (around −200 °C). Because of its low cost, liquid nitrogen is also often used when such low temperatures are not strictly necessary, such as refrigeration of food, livestock, freezing pipes to halt flow when valves are not present, and consolidating unstable soil by freezing whenever excavation is going on underneath. Safety Gas Although nitrogen is non-toxic, when released into an enclosed space it can displace oxygen, and therefore presents an hazard. This may happen with few warning symptoms, since the human is a relatively poor and slow low-oxygen (hypoxia) sensing system. An example occurred shortly before the launch of the first Space Shuttle mission in 1981, when two technicians died from asphyxiation after they walked into a space located in the Shuttle's that was pressurised with pure nitrogen as a precaution against fire. When inhaled at high (more than about 4 bar, encountered at depths below about 30 m in ), nitrogen is an anesthetic agent, causing, a temporary state of mental impairment similar to intoxication.
Nitrogen dissolves in the and body fats. Rapid decompression (as when divers ascend too quickly or astronauts decompress too quickly from cabin pressure to spacesuit pressure) can lead to a potentially fatal condition called (formerly known as caisson sickness or the bends), when nitrogen bubbles form in the bloodstream, nerves, joints, and other sensitive or vital areas.
Bubbles from other 'inert' gases (gases other than carbon dioxide and oxygen) cause the same effects, so replacement of nitrogen in may prevent nitrogen narcosis, but does not prevent decompression sickness. Liquid As a liquid, liquid nitrogen can be dangerous by causing on contact, although the provides protection for very short exposure (about one second). Ingestion of liquid nitrogen can cause severe internal damage.
For example, in 2012, a young woman in England had to have her stomach removed after ingesting a cocktail made with liquid nitrogen. Because the liquid-to-gas of nitrogen is 1:694 at 20 °C, a tremendous amount of force can be generated if liquid nitrogen is rapidly vaporised in an enclosed space. In an incident on January 12, 2006 at, the pressure-relief devices of a tank of liquid nitrogen were malfunctioning and later sealed. As a result of the subsequent pressure buildup, the tank failed catastrophically. The force of the explosion was sufficient to propel the tank through the ceiling immediately above it, shatter a reinforced concrete beam immediately below it, and blow the walls of the laboratory 0.1–0.2 m off their foundations. Liquid nitrogen readily evaporates to form gaseous nitrogen, and hence the precautions associated with gaseous nitrogen also apply to liquid nitrogen. For example, are sometimes used as a safety precaution when working with liquid nitrogen to alert workers of gas spills into a confined space.
Vessels containing liquid nitrogen can from air. The liquid in such a vessel becomes increasingly enriched in oxygen (boiling point −183 °C) as the nitrogen evaporates, and can cause violent oxidation of organic material. See also. References.
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