Nitrous+dioxide

Nitrogen dioxide is the chemical compound with the formula NO 2. It is one of several nitrogen oxides. NO 2 is an intermediate in the industrial synthesis of nitric acid, millions of tons of which are produced each year. This reddish-brown toxic gas has a characteristic sharp, biting odor and is a prominent air pollutant. Nitrogen dioxide is a paramagnetic bent molecule. Nitrogen dioxide plays an important role in the process of forming ozone. It's one of the gases which cause the acid rain. Nitrogen dioxide can be dissolved in water and produce nitric acid and nitric oxide. It does harm to our respiratory tract, lung and skin. In lab people always use copper and nitric acid (high concentration) to produce nitrogen dioxide and use sodium hydroxide to eliminate the remaining nitrogen dioxide. It will be converted into N 2 O 4 easily in a low temperature.

Molecular Formula: NO 2 Molar Mass: 46.0055 g/mol



Workers are exposed to combustion-produced NO 2 in various occupations, including arc welders, firefighters, military and aerospace personnel, and those working with explosives. Nitric oxide (NO), NO 2 and other oxides of nitrogen are formed from nitrogen and oxygen during high-temperature combustion. NO is oxidized to NO 2, a precursor of ozone (O 3 ). NO, NO 2, and nitrogen tetroxide (N 2 O 4 ) almost always occur together; hence, the terms oxides of nitrogen and nitrogen oxides are used in literature to refer to these molecules. The term NO x is used most often in air pollution literature in reference to the oxides of nitrogen.

NO 2 toxicity is also observed in environments where NO 2 is formed from noncombustion sources. These include silo fillers, where nitrogen oxides are a byproduct of anaerobic fermentation of crops, and indoor ice skating rinks, [|[1]] where the gas is generated by the propane-driven ice cleaning machine, the Zamboni. In addition, NO 2 from automobile exhaust smokestack emissions are thought to be a major contributor to the toxic effects of air pollution.

NO 2 is a deep lung irritant that can produce pulmonary edema and fatality if inhaled at high concentrations. The effects of NO 2 depend on the level and duration of exposure. Exposure to moderate NO 2 levels (50 ppm) may produce cough, hemoptysis, dyspnea, and chest pain. Exposure to higher concentrations of NO 2 (>100 ppm) can produce pulmonary edema that may be fatal or may lead to bronchiolitis obliterans.

Molecular properties
Nitrogen dioxide has a [|molar mass] of 46.0055, which makes it heavier than air, whose average molar mass is 28.8. According to the [|ideal gas law], NO 2 is therefore more dense than air. The [|bond] length between the nitrogen atom and the oxygen atom is 119.7 [|pm]. This bond length is consistent with a bond order of one and a quarter, as in [|ozone] (O 3 ). The [|ground] [|electronic state] of nitrogen dioxide is a [|doublet state], since there is one unpaired bonding electron delocalised over both bonds, hence the bond order of one and a quarter.

Occurrence
NO 2 exists in equilibrium with the colorless gas [|dinitrogen tetroxide] ( N 2 O 4 ): 2 NO 2 N 2 O 4 The equilibrium is characterized by ΔH = −57.23 kJ/mol, which is exothermic. Resulting from an [|endergonic] reaction at higher temperatures, the paramagnetic monomer is favored. Colorless diamagnetic N 2 O 4 can be obtained as a solid melting at m.p. −11.2 °C. [|[3]]

Preparation and reactions
Nitrogen dioxide typically arises via the oxidation of [|nitric oxide] by oxygen in air: [|[3]] 2 NO + O 2 → 2 NO 2 In the laboratory, NO2 can be prepared in a two step procedure by thermal decomposition of [|dinitrogen pentoxide], which is obtained by dehydration of nitric acid: 2 HNO 3 → N 2 O 5 + H 2 O 2  N 2 O 5 → 4 NO 2 + O 2 The thermal decomposition of some metal nitrates also affords NO 2 : 2 Pb(NO 3 ) 2 → 2 PbO + 4 NO 2 + O 2 Alternatively, reduction of concentrated nitric acid by metal (such as copper). 4 HNO 3 + Cu → Cu(NO 3 ) 2 + 2 NO 2 +2 H 2 O Or finally by adding concentrated nitric acid over tin. [|Stannic acid] is produced as byproduct. 4HNO 3 + Sn → H 2 O + H 2 SnO 3 + 4 NO 2

Main reactions
The chemistry of nitrogen dioxide has been investigated extensively. At 150 °C, NO 2 decomposes with release of oxygen via an endothermic process (ΔH = 114 kJ/mol): 2 NO 2 → 2 NO + O 2 As suggested by the weakness of the N–O bond, NO 2 is a good oxidizer. Consequently, it will combust, sometimes explosively, with many compounds, such as [|hydrocarbons]. It hydrolyzes to give [|nitric acid] and [|nitrous acid] : 2 NO 2 / N 2 O 4 + H 2 O → HNO 2 + HNO 3 This reaction is one step in the [|Ostwald process] for the industrial production of nitric acid from ammonia. [|[4]] Nitric acid decomposes slowly to nitrogen dioxide, which confers the characteristic yellow color of most samples of this acid: 4 HNO 3 → 4 NO 2 + 2 H 2 O + O 2 NO 2 is used to generate anhydrous metal nitrates from the oxides: [|[3]] MO + 3 NO 2 → 2 M(NO 3 ) 2 + NO Alkyl and metal iodides give the corresponding nitrites: 2 CH 3 I + 2 NO 2 → 2 CH 3 NO 2 + I 2 TiI 4 + 4 NO 2 → Ti(NO 2 ) 4 + 2 I 2