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| Names | |
|---|---|
| IUPAC name Dinitrogen pentoxide | |
| Other names Nitric anhydride Nitronium nitrate Nitryl nitrate DNPO Anhydrous nitric acid | |
| Identifiers | |
3D model (JSmol) |
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| ChEBI | |
| ChemSpider |
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| ECHA InfoCard | 100.030.227 |
| EC Number |
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| UNII | |
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| Properties | |
| N2O5 | |
| Molar mass | 108.01 g/mol |
| Appearance | white solid |
| Density | 2.0 g/cm3[1] |
| Boiling point | 33 °C (91 °F; 306 K) sublimes[1] |
| reacts to giveHNO3 | |
| Solubility | soluble inchloroform negligible inCCl4 |
| −35.6×10−6 cm3 mol−1 (aq) | |
| 1.39 D | |
| Structure[2] | |
| Hexagonal,hP14 | |
| P63/mmc No. 194 | |
a = 0.54019 nm,c = 0.65268 nm | |
Formula units (Z) | 2 |
| planar,C2v (approx.D2h) N–O–N ≈ 180° | |
| Thermochemistry[3] | |
| 143.1 J K−1 mol−1 (s) 95.3 J K−1 mol−1 (g) | |
Std molar entropy(S⦵298) | 178.2 J K−1 mol−1 (s) 355.7 J K−1 mol−1 (g) |
Std enthalpy of formation(ΔfH⦵298) | −43.1 kJ/mol (s) +13.3 kJ/mol (g) |
Gibbs free energy(ΔfG⦵) | 113.9 kJ/mol (s) +117.1 kJ/mol (g) |
| Hazards | |
| Occupational safety and health (OHS/OSH): | |
Main hazards | strong oxidizer, forms strong acid in contact with water |
| NFPA 704 (fire diamond) | |
| Flash point | Non-flammable |
| Related compounds | |
| Nitrous oxide Nitric oxide Dinitrogen trioxide Nitrogen dioxide Dinitrogen tetroxide | |
Related compounds | Nitric acid |
Except where otherwise noted, data are given for materials in theirstandard state (at 25 °C [77 °F], 100 kPa). | |
Dinitrogen pentoxide (also known asnitrogen pentoxide ornitric anhydride) is thechemical compound with theformulaN2O5. It is one of the binarynitrogen oxides, a family of compounds that contain onlynitrogen andoxygen. It exists as colourless crystals that sublime slightly above room temperature, yielding a colorless gas.[4]
Dinitrogen pentoxide is an unstable and potentially dangerous oxidizer that once was used as areagent when dissolved inchloroform fornitrations but has largely been superseded bynitronium tetrafluoroborate (NO2BF4).
N2O5 is a rare example of a compound that adopts two structures depending on the conditions. The solid is a salt,nitronium nitrate, consisting of separatenitronium cations[NO2]+ andnitrate anions[NO3]−; but in the gas phase and under some other conditions it is acovalently-bound molecule.[5]
N2O5 was first reported by the French chemistHenri Deville in 1840, who prepared it by treatingsilver nitrate (AgNO3) withchlorine.[6][7]
Pure solidN2O5 is asalt, consisting of separated linearnitronium ionsNO+2 and planar trigonalnitrate anionsNO−3. Bothnitrogen centers haveoxidation state +5. It crystallizes in the space groupD4
6h (C6/mmc) withZ = 2, with theNO−3 anions in theD3h sites and theNO+2 cations inD3d sites.[8]
The vapor pressureP (in atm) as a function of temperatureT (inkelvin), in the range 211 to 305 K (−62 to 32 °C), is well approximated by the formula
being about 48 torr at 0 °C, 424 torr at 25 °C, and 760 torr at 32 °C (9 °C below the melting point).[9]
In the gas phase, or when dissolved in nonpolarsolvents such ascarbon tetrachloride, the compound exists ascovalently-bonded moleculesO2N−O−NO2. In the gas phase, theoretical calculations for the minimum-energy configuration indicate that theO−N−O angle in each−NO2 wing is about 134° and theN−O−N angle is about 112°. In that configuration, the two−NO2 groups are rotated about 35° around the bonds to the central oxygen, away from theN−O−N plane. The molecule thus has a propeller shape, with one axis of 180° rotational symmetry (C2)[10]
When gaseousN2O5 is cooled rapidly ("quenched"), one can obtain themetastable molecular form, which exothermically converts to the ionic form above −70 °C.[11]
GaseousN2O5 absorbsultraviolet light with dissociation into thefree radicalsnitrogen dioxideNO2• andnitrogen trioxideNO3• (uncharged nitrate). The absorption spectrum has a broad band with maximum at wavelength 160 nm.[12]
A recommended laboratory synthesis entails dehydratingnitric acid (HNO3) withphosphorus(V) oxide:[11]
Another laboratory process is the reaction oflithium nitrateLiNO3 andbromine pentafluorideBrF5, in the ratio exceeding 3:1. The reaction first formsnitryl fluorideFNO2 that reacts further with the lithium nitrate:[8]
The compound can also be created in the gas phase by reactingnitrogen dioxideNO2 orN2O4 withozone:[13]
However, the productcatalyzes the rapid decomposition of ozone:[13]
Dinitrogen pentoxide is also formed when a mixture of oxygen and nitrogen is passed through an electricdischarge.[8] Another route is the reactions ofPhosphoryl chloridePOCl3 ornitryl chlorideNO2Cl withsilver nitrateAgNO3[8][14]
Dinitrogen pentoxide reacts with water (hydrolyses) to producenitric acidHNO3. Thus, dinitrogen pentoxide is theanhydride of nitric acid:[11]
Solutions of dinitrogen pentoxide in nitric acid can be seen as nitric acid with more than 100% concentration. The phase diagram of the systemH2O−N2O5 shows the well-known negativeazeotrope at 60%N2O5 (that is, 70%HNO3), a positive azeotrope at 85.7%N2O5 (100%HNO3), and another negative one at 87.5%N2O5 ("102%HNO3").[15]
The reaction withhydrogen chlorideHCl also gives nitric acid andnitryl chlorideNO2Cl:[16]
Dinitrogen pentoxide eventually decomposes at room temperature intoNO2 andO2.[17][13] Decomposition is negligible if the solid is kept at 0 °C, in suitably inert containers.[8]
Dinitrogen pentoxide reacts withammoniaNH3 to give several products, includingnitrous oxideN2O,ammonium nitrateNH4NO3,nitramideNH2NO2 andammonium dinitramideNH4N(NO2)2, depending on reaction conditions.[18]
Dinitrogen pentoxide between high temperatures of 600 and 1,100 K (327–827 °C), is decomposed in two successive stoichiometric steps:
In the shock wave,N2O5 has decomposed stoichiometrically intonitrogen dioxide andoxygen. At temperatures of 600 K and higher, nitrogen dioxide is unstable with respect tonitrogen oxideNO and oxygen. The thermal decomposition of 0.1 mM nitrogen dioxide at 1000 K is known to require about two seconds.[19]
Apart from the decomposition ofN2O5 at high temperatures, it can also be decomposed incarbon tetrachlorideCCl4 at 30 °C (303 K).[20] BothN2O5 andNO2 are soluble inCCl4 and remain in solution while oxygen is insoluble and escapes. The volume of the oxygen formed in the reaction can be measured in a gas burette. After this step we can proceed with the decomposition, measuring the quantity ofO2 that is produced over time because the only form to obtainO2 is with theN2O5 decomposition. The equation below refers to the decomposition ofN2O5 inCCl4:
And this reaction follows the first orderrate law that says:
N2O5 can also be decomposed in the presence ofnitric oxideNO:
The rate of the initial reaction between dinitrogen pentoxide and nitric oxide of the elementary unimolecular decomposition.[21]
Dinitrogen pentoxide, for example as a solution inchloroform, has been used as a reagent to introduce the−NO2 functionality inorganic compounds. Thisnitration reaction is represented as follows:
where Ar represents anarene moiety.[22] The reactivity of theNO+2 can be further enhanced with strong acids that generate the "super-electrophile"HNO2+2.
In this use,N2O5 has been largely replaced bynitronium tetrafluoroborate[NO2]+[BF4]−. This salt retains the high reactivity ofNO+2, but it is thermally stable, decomposing at about 180 °C (intoNO2F andBF3).
Dinitrogen pentoxide is relevant to the preparation of explosives.[7][23]
In theatmosphere, dinitrogen pentoxide is an important reservoir of theNOx species that are responsible forozone depletion: its formation provides anull cycle with whichNO andNO2 are temporarily held in an unreactive state.[24]Mixing ratios of several parts per billion by volume have been observed in polluted regions of the nighttime troposphere.[25] Dinitrogen pentoxide has also been observed in the stratosphere[26] at similar levels, the reservoir formation having been postulated in considering the puzzling observations of a sudden drop in stratosphericNO2 levels above 50 °N, the so-called 'Noxon cliff'.
Variations inN2O5 reactivity inaerosols can result in significant losses in troposphericozone,hydroxyl radicals, andNOx concentrations.[27] Two important reactions ofN2O5 in atmospheric aerosols are hydrolysis to formnitric acid[28] and reaction withhalide ions, particularlyCl−, to formClNO2 molecules which may serve as precursors to reactive chlorine atoms in the atmosphere.[29][30]
N2O5 is a strong oxidizer that forms explosive mixtures with organic compounds andammonium salts. The decomposition of dinitrogen pentoxide produces the highly toxicnitrogen dioxide gas.
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