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The chemical formula is NO2. It is one of the important nitrogen oxides. Nitrogen dioxide can be prepared by the oxidation of nitric oxide. In the laboratory, small amounts of NO2 are produced by reacting copper with concentrated nitric acid, or by heating heavy metal nitrates: 2Pb(NO3)2 → 2PbO + 4NO2 + O2↑. The nitrogen dioxide molecule has an angular shape, indicating that the nitrogen atom is sp2-hybridized. Its bond length (120 pm) lies between that of a single bond (143 pm) and a double bond (118 pm), suggesting the presence of delocalized π bonds (π33) in addition to localized σ bonds; similar to NO, it is paramagnetic (with a total of 17 valence electrons). NO2 is a reddish-brown, toxic gas; its molecule is polar (μ=0.29D). It has a higher solubility in water than NO, and when it reacts with cold water it produces nitric acid and nitrous acid. Therefore, NO2 is also known as a mixed anhydride: 2NO2 + H2O = HNO3 + HNO2. When dissolved in hot water, NO2 forms nitric acid and NO: 3NO2 + H2O = 2HNO3 + NO. The melting point of NO2 is -11°C, while its boiling point is 21°C. At low temperatures, it tends to polymerize into colorless, diamagnetic dinitrogen tetroxide, N2O4. Pure N2O4 exists only in the solid state; upon melting it becomes a yellow liquid that contains about 1% NO2. At 100°C, it contains around 90% NO2, and above 150°C it decomposes into NO and O2. The N2O4 molecule has a planar structure, with a 68-membered π bond perpendicular to the plane. NO2 has strong oxidizing properties; substances such as carbon, sulfur, and phosphorus can continue to burn in its presence. Potassium metal catches fire immediately upon contact with NO2, and NO2 can oxidize SO2 to SO3. Based on the value of the standard electrode potential (E), the oxidizing power of NO2 is similar to that of bromine, and it is stronger than that of nitric acid. The strong oxidizing power of concentrated nitric acid is also related to its high content of NO2; for example, colorless nitric acid without NO2 initially finds it difficult to react with Cu, but once NO2 is produced, the reaction proceeds rapidly. NO2 acts as a catalyst for the transfer of electrons between the reducing agent and nitric acid. NO2 has very weak reducibility and is oxidized only in the presence of very strong oxidants. For example, MnO4+ + NO2 + 6H+ = Mn2+ + NO3- + 3H2O. NO2 can react violently with many organic compounds, and it is used as a catalyst in the nitroprocess for the production of sulfuric acid.