Oxide of nitrogen

Nitrogen reacts with oxygen to form five oxides, as shown in the table on the adjoining page. The two most important with regard to air pollution are NO and NO2, which are sometimes referred to together as NO. In some circumstances, the formula may also include other oxides of nitrogen. 

Nitrogen and oxygen coexist in the atmosphere without reacting to any significant extent. At room temperature, for example, air normally consists of no more than about 0.000,000,000, 1 (1 x 10 °) ppm of NO, formed in the following reaction  

The reaction rate for this reaction accelerates rapidly, however, with increases in temperature. For example, at about 1,000°C, the concentration of NO reaches about 100 ppm and at about 1,500°C, it reaches nearly 1,000 ppm. 

The most common source of nitrogen oxides, therefore, is high-temperature combustion processes, such as those that take place in automobiles and trucks, in electrical power generating plants, and in industrial processes. Residential sources, such as gas stoves and home heaters, are also responsible for the release of significant amounts of NO into the atmosphere. At the end of the 20th century, the EPA reported that motor vehicles were responsible for 49 percent of all NO released into the atmosphere in the United States  

NO2 nitrogen dioxide also occurs as N2O4, dinitrogen tetroxide 

Nitrogen monoxide (NO) and nitrogen dioxide (NO2) are almost unique in being small, common molecules which contain unpaired electrons. 

The fraction of the total that is emitted as NO clearly depends on the conditions associated with the specific combustion process. While most (typically 90%) of the NOx emitted is believed to be in the form of NO, the fraction of N02 can vary from less than 1% to more than 30% (e.g., Lenner, 1987). 

FIGURE 2.2 Contribution of various sources to total anthropogenic NOx emissions in the United States in 1996 (from EPA, 1999). 

FIGURE 2.3 NOx emissions in million tons of equivalent N02 for the period 1970 to 1986 for Asia, Europe, North America, and the USSR (from Hameed and Dignon, 1992). 

There is also some NO produced from the oxidation of NH3 by photochemical processes in oceans and by some terrestrial plants (e.g., Wildt et al., 1997). 

Nitrous oxide (N20, laughing gas ) is also produced by biological processes and, to a lesser extent, by anthropogenic processes (see Chapter 14.B.2c). While 

The physiological effects of N2O (laughing gas, anaesthetic) and NO2 (acrid, corrosive fumes) have been known from the earliest days, and the environmental problems of NOj from automobile exhaust fumes and as a component in photochemical smog are well known in all industrial countries. 

NO is now recognized as a key neuro transmitter in humans and other animals and its biologically triggered synthesis is implicated in cardiovascular pharmacology, hypertension, impotence, immunology and other vital functions.NO and NO2 are important in 

Lee (ed.). Nitrogen Oxides and their Effects on Health, Ann Arbor Publishers, Michigan, 1980, 382 pp. 

Bosch and F. J. J. Janssen, Catalytic Reduction of Nitrogen Oxides, Elsevier, Amsterdam, 1988, 164 pp. 

The oxides of nitrogen played an important role in exemplifying Dalton s law of multiple proportions which led up to the formulation of his atomic theory (1803-8), and they still pose some fascinating problems in bonding theory. Their formulae, molecular structure, and physical appearance are briefly summarized in Table 11.7 and each compound is discussed in turn in the following sections. 

Therefore the equation balances in terms of oxidation state 

In reaction 15.56, which elements are oxidized and which reduced Confirm that the reaction balances in terms of changes in oxidation states. [Ans. N, oxidized F, reduced] 

Which elements undergo redox changes in reaction 15.59 Confirm that the equation balances in terms of the oxidation state changes. [Ans. N, reduced half of the Cl, oxidized] 

Are reactions 15.71,15.72 and 15.73 redox reactions Confirm your answer by determining the oxidation states of the N atoms in the reactants and products in each equation. 

Confirm that reaction 15.87 is a redox process, and that the equation balances with respect to changes in oxidation states for the appropriate elements. 

As in group 14, the first element of group 15 stands apart in forming oxides in which (/ -/ )7r-bonding is important. Table 14.6 lists selected properties of nitrogen oxides, excluding NO3 which is an unstable radical NO2 exists in equilibrium with N2O4. 

Dinitrogen monoxide (Table 14.6) is usually prepared by decomposition of solid ammonium nitrate (equation 14.87, compare reaction 14.6) but the aqueous solution reaction 14.88 is useful for obtaining a purer product. For further detail on the oxidation of NH2OH to N2O, see Section 14.5. 

Dinitrogen monoxide has a faint, sweet odour. It dissolves in water to give a neutral solution, but does not react to any significant extent. The position of equilibrium 14.89 is far to the left. 

Dinitrogen monoxide (Table 15.6) is usually prepared by decomposition of solid ammonium nitrate (eq. 15.93, 

Name Dinitrogen monoxide Nitrogen monoxide Dinitrogen trioxide Nitrogen dioxide Dinitrogen tetraoxide Dinitrogen pentaoxide 

Physical appearance Colourless gas Colourless gas Blue solid or liquid Brown gas Colourless solid or liquid, but see text Colourless solid, stable below 273 K 

Magnetic properties Diamagnetic Paramagnetic Diamagnetic Paramagnetic Diamagnetic Diamagnetic 

In the late 1960s, University of California physical chemist Harold S. Johnston (1920-) and Dutch atmospheric chemist Paul J. Crutzen (1933-) independently proposed that emissions of nitrogen oxides from supersonic transport aircraft (SST) flying through the stratosphere would harm the stratosphere s protective layer of ozone. 

The effect of a fleet of SSTs would be deleterious to the stratosphere s ozone. What eventually killed the SST program, though, was economics—in the mid-1970s, fuel became too expensive. 

The chemical inertness of CFCs posed a problem, however. Refrigeration and air-conditioning systems sometimes leak. Disused refrigerators tend to end up in landfills. 

The mechanism Rowland and Molina proposed is completely analogous to the mechanism for NO molecules  

Repeat these reactions tens of thousands of times. 

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Nitrogen monoxide is the most stable of all the oxides of nitrogen. It can be prepared in small amounts by direct combination of the... 

With more concentrated nitric acid, oxides of nitrogen are formed. 

Equip a 500 ml. three necked flask with a reflux condenser, a mercury-sealed mechanical stirrer and separator funnel, and support it on a water bath. Attach an absorption device (Fig. II, 8, 1, c) to the top of the condenser (1). Place 134 g. (152 ml.) of A.R, benzene and 127 g. of iodine in the flask, and heat the water bath to about 50° add 92 ml. of fuming nitric acid, sp. gr. 1-50, slowly from the separatory funnel during 30 minutes. Oxides of nitrogen are evolved in quantity. The temperature rises slowly without the application of heat until the mixture boils gently. When all the nitric acid has been introduced, reflux the mixture gently for 15 minutes. If iodine is still present, add more nitric acid to the warm solution until the purple colour (due to iodine) changes to brownish-red. 

It is advisable to add the sodium nitrite solution, particularly in preparations on a larger scale, through a separatory or dropping funnel with the tip of the stem extending well below the sui-face of the liquid tliis will prevent loss of nitrous acid by surface decomposition into oxides of nitrogen. 

Method 1. Place 20 g. of crude benzoin (preceding Section) and 100 ml. of concentrated nitric acid in a 250 ml. round-bottomed flask. Heat on a boiling water bath (in the fume cupboard) with occasional shaking until the evolution of oxides of nitrogen has ceased (about 1 -5 hours). Pour the reaction mixture into 300-400 ml. of cold water contained in a beaker, stir well until the oil crystallises completely as a yellow solid. Filter the crude benzil at the pump, and wash it thoroughly with water to remove the nitric acid. RecrystaUise from alcohol or methylated spirit (about 2-5 ml. per gram). The yield of pure benzil, m.p. 94-96°, is 19 g. 

Dissolve 1.000 g Au in 10 ml of hot HNO3 by dropwise addition of HCI, boil to expel oxides of nitrogen and chlorine, and dilute to volume. Store in amber container away from light. 

Fig. 7. NO formation for the Provo-Orem bus mn at a compression ratio of 12 1 at 30°C, 3000 rpm, where A is brake mean effective pressure B, brake thermal efficiency and C, oxides of nitrogen, (a) Effect of equivalence ratio, ( ), at a water/H2 mass ratio of 6.0 and spark = 17° before top-dead (BTC) and (b), effect of water injection where (j) = 0.60 and spark = 14°BTC. To convert MPa to psi, multiply by 14.

Oxides of nitrogen, NO, can also form. These are generally at low levels and too low an oxidation state to consider water scmbbing. A basic reagent picks up the NO2, but not the lower oxidation states the principal oxide is usually NO, not NO2. Generally, control of NO is achieved by control of the combustion process to minimize NO, ie, avoidance of high temperatures in combination with high oxidant concentrations, and if abatement is required, various approaches specific to NO have been employed. Examples are NH injection and catalytic abatement (43). 

Air Pollution. Particulates and sulfur dioxide emissions from commercial oil shale operations would require proper control technology. Compliance monitoring carried out at the Unocal Parachute Creek Project for respirable particulates, oxides of nitrogen, and sulfur dioxide from 1986 to 1990 indicate a +99% reduction in sulfur emissions at the retort and shale oil upgrading faciUties. No violations for unauthorized air emissions were issued by the U.S. Environmental Protection Agency during this time (62). 

Unpiotonated hydioxylamine is oxidized rapidly by ozone, / = 2.1 X 10 (39). The reaction of ozone with the lower oxides of nitrogen (NO and NO2) is also rapid and quantitative the end product is nitrogen pentoxide, which is also a catalyst for the decomposition of ozone (45). Nitrous oxide, however, reacts slowly (k < 10 ) (39). Nitrogen-containing anions, eg, nitrite and cyanide, also ate oxidized by ozone (39). Nitrite is oxidized to nitrate (fc = 3.7 X 10 and cyanide is oxidized rapidly to cyanate (fc = 2.6 X 10 (46) and 10 -10 (39)). Cyanate, however, is oxidized slowly. 

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