Control of Nitrogen Oxides

1. Control of Nitrogen Oxides 朱信 Hsin Chu Professor Dept. of Environmental Engineering National Cheng Kung University

2. 1. An Overview of the Nitrogen Oxides Problem • Most of the world’s nitrogen is in the atmosphere as an inert gas. • In crustal rocks it is the 34th most abundant element with an abundance of only ≈ 20 ppm.

3. Although nitrogen forms eight different oxides, our principal air pollution interest is in the two most common oxides, nitric oxide (NO) and nitrogen dioxide (NO2). • In addition, we are beginning to be concerned with nitrous oxide (N2O). It may be a significant contributor to global warming and to the possible destruction of the ozone layer.

4. 1.1 Comparison with Sulfur Oxides • Fig. 12.1 (next slide) shows part of how nitrogen moves in the environment, as a result of human activities.

6. Nitrogen oxides are often lumped with sulfur oxides as air pollution control problems because of the similarities between the two:(1) Nitrogen oxides and sulfur oxides react with water and oxygen in the atmosphere to form nitric sulfuric acids, respectively. These two acids are the principal contributors to acid rain.

7. (2) Both undergo atmospheric transformations leading to or contributing to the formation of PM10 and PM2.5 in urban areas.(3) Both are released into the atmosphere in large quantities.(4) Both are released to the atmosphere by large combustion sources, particularly coal combustion sources. Table 12.1 (next slide) shows the emission sources of NOX in the U.S. in 1997.

9. We see that vehicles contribute almost half of the total. • Natural gas appears in this table, but not in the SO2 table. • The largest contributors to the “All other sources” are nontransportation internal combustion engines and residential combustion.

10. There are also major differences between nitrogen oxides and sulfur oxides: (1) Motor vehicles are the major emitter of nitrogen oxides, but a very minor source of sulfur oxides. (2) Sulfur oxides are formed from the sulfur contaminants in fuels or the unwanted sulfur in sulfide ores. Although some of nitrogen oxides emitted to the atmosphere are due to nitrogen contaminants in fuels, most are not.

11. Most of nitrogen oxides are formed by the reaction of atmospheric nitrogen with oxygen in high-temperature flames. (3) The formation of nitrogen oxides in flames can be greatly reduced by manipulating the time, temperature, and oxygen content of the flames. No such reductions are possible with sulfur oxides.

12. (4) The ultimate fate of sulfur oxides removed in pollution control or fuel-cleaning process is to be turned into CaSO4•2H2O, which is an innocuous, low-solubility solid, and to be placed in landfills. • There is no correspondingly cheap, innocuous, and insoluble salt of nitric acid, so landfilling is not a suitable fate for the nitrogen oxides.

13. The ultimate fate of nitrogen oxides that we wish to keep out of the atmosphere is to be converted to gaseous nitrogen and oxygen and be returned to the atmosphere. (5) It is relatively easy to remove SO2 from combustion gases by dissolving SO2 in water and reacting it with alkali. • Collecting nitrogen oxides is not nearly as easy this way because NO, the principal nitrogen oxide present in combustion gas streams, has a very low solubility in water.

14. NO must undergo a two-step process to form an acid: NO + 0.5 O2↔ NO2 (1) 3 NO2 +H2O → 2HNO3 + NO (2)The first reaction is relatively slow. • It is fast enough in the atmosphere to lead to the formation of acid precipitation in the several hours or days that the polluted air travels before encountering precipitation, but slow-enough that it does not remove significant quantities of NO in the few seconds that a contaminated gas spends in a wet limestone scrubber used for SO2 control.

15. 1.2 Reactions in the Atmosphere • NO is a colorless gas that has some harmful effects on health, but these effects are substantially less than those of an equivalent amount of NO2. • NO2, a brown gas, is a serious respiratory irritant.

16. Our principal concern with NOX is that nitrogen oxides contribute to the formation of ozone, O3, which is a strong respiratory irritant and one of the principal constituents of urban summer eye-and nose-irritating smog. • The overall reaction is: NO + HC + O2 + sunlight → NO2 + O3

17. NO, emitted during the morning commuter rush, is oxidized in the atmosphere to NO2 over a period of several hours. • The NO2 thus formed then reacts with HC to form O3. • The O3 peak occurs after the NO2 peak.

18. 1.3 NO and NO2 Equilibrium • The most important reactions for producing NO and NO2 in flames are: N2 + O2↔ 2NO (3)and Eq. (1). • For any chemical reaction at equilibrium, the Gibbs free energy is at a minimum for the reaction’s temperature and pressure.

19. From that condition it follows that where ∆Go = standard Gibbs free energy change K = equilibrium constant R = the universal constant T = absolute temperature (K or oR) • Using the published Go values for the reactions in Eqs. (1) and (3) we may construct Table 12.2 (next slide).

21. Here we show Kp, the equilibrium constant based on taking the standard states as perfect gases at 1 atm pressure, and expressing gas concentrations as partial pressures. • Because the number of mols does not change in Eq. (3), its Kp is dimensionless. • However, in Eq. (1) the number of mols decreases by ½, so that Kp has the dimension (atm)-0.5.

22. Example1 • Calculate the equilibrium concentrations of NO and NO2 for air that is held at 2000 K = 3140oF long enough to reach equilibrium. • Assume that the only reactions of interest are Eqs. (1) and (3).

23. Solution:The definitions of the two equilibrium constants are: • Solving Eq. (3) for the equilibrium concentration of NO, we find

24. Substituting values, including the value of K3 from Table 12.2, into this equation, we see • Solving Eq. (6) for [NO2], we find

25. Substituting values into this equation, we see • Following the same procedure as in Example 1, we may make up Table 12.3 (next slide), which shows the calculated equilibrium concentrations of NO and NO2 at various temperatures.

27. The starting gas that is 78% nitrogen, 4% oxygen is more representative of combustion gases in which most of the oxygen has been consumed by the combustion. • The values from Table 12.3 are shown in Figs. 12.2 (next slide) and 12.3 (second slide).

30. From Table 12.3 and Figs. 12.2 and 12.3 we see:(1) If the atmosphere were at equilibrium (at a temperature near 300 K = 80oF), it would have less than a part per billion of NO or NO2. • The concentrations of NO and NO2 observed in cities in the world often exceed these equilibrium values, so that equilibrium alone is not a satisfactory guide to the presence of NO and NO2 in the atmosphere.

31. (2) The equilibrium concentration of NO increases dramatically with increasing temperature. The rapid increase begins at about 2000~2500oF (1367-1644 K). (3) At low temperatures the equilibrium concentration of NO2 is much higher than that of NO, whereas at high temperatures the reverse is true.

32. (4) We are likely to get to the temperatures where NO and NO2 are formed from atmospheric N2 and O2 only in flames and in lightning strikes. • Lightning strikes are a major global source of NOX, but combustion in our vehicles and factories is the main source of NOX in heavily populated areas.

33. 1.4 Thermal, Prompt, and Fuel NOX • Thermal nitrogen oxides, which are generally the most significant, are formed by the simple heating of oxygen and nitrogen, either in a flame or by some other external heating. • Prompt refers to the nitrogen oxides that form very quickly as a result of the interaction of nitrogen and oxygen with some of the active hydrocarbon species derived from the fuel in the fuel-rich parts of flames.

34. Prompt NOX are not observed in flames of fuels with no carbon, e.g., H2. • Fuel nitrogen oxides is formed by conversion of some of the nitrogen originally present in the fuel to NOX.

35. Coal and some high-boiling petroleum fuels contain significant amounts of organic nitrogen; low-boiling petroleum fuels and natural gas contain practically none. • Fig. 12.4 (next slide) shows estimates of the contribution of the thermal, fuel, and prompt mechanisms to the NOX emissions from coal combustion.

37. 2. Thermal NO • From Fig. 12.4, the thermal mechanism is the most important of the three ways of making NO at the highest temperatures. 2.1 The Zeldovich Kinetics of Thermal NO Formation • Example 1 showed that a specific equilibrium concentration of NO could be reached at certain temperature. • Here we inquire how rapidly the mixture approaches that equilibrium value.

38. The reactions shown in Eqs. (1) and (3) do not proceed as written in those equations. • Rather, they proceed by means of intermediate steps involving highly energetic particles called free radicals. • The free radicals most often involved in combustion reactions are O, N, OH, H, and hydrocarbons that have lost one or more hydrogens, e. g., CH3 or CH2.

39. These materials are very reactive and energetic and can exist in significant concentrations only at high temperatures. • In principle they can be formed by equilibrium reactions like the following: N2↔ 2N O2 ↔ 2O H2O ↔ H + OH

40. However, the reactions are conventionally written with an M on both sides of the equation to indicate that another molecule, which is not chemically changed in the reaction, must collide with the N2 or O2 molecule to supply or remove energy for the reaction to occur, e.g., N2 + M ↔ 2N +M

41. The concentration of M, which can be any other gas molecule (e.g., O2 or N2 or H2O) does not influence the equilibrium but does influence the rate of the reactions. • The most widely quoted mechanism for the thermal NO formation reaction is that of Zeldovich. • It assumes that O radicals attack N2 molecules by O + N2↔ NO + Nand that N radicals can form NO by N + O2 ↔ NO + O

42. If one assumes that the O radicals are in equilibrium with O2, that the concentration of N radicals is not changing significantly with time, and that one term is small compared with the others, one can simplify the resulting kinetic equations to:where kf and kb are the forward and backward reaction rate constants.

43. Then one can observe that the equilibrium value of the NO concentration, [NO]e, is given by Eq. (5), and that the equilibrium constant is related to the two rate constants by • One may show this by setting the rate of change of concentration in Eq. (7) to zero and comparing the result to Eq. (5).

44. Making this substitution and rearranging in Eq. (7), one finds:which may be integrated and rearranged (with an assumption of zero NO at time zero) to where

45. Example 2 • Estimate the concentration of NO in a sample containing 78% N2 and 4% O2 that is hold for one second at 2000 K = 3140oF, according to the Zeldovich thermal mechanism. • Solution:From Table 12.3, we know the equilibrium concentration, [NO]e, is 3530 ppm.

46. Seinfeld suggests a value for kb of 4.1  1013 exp (-91600/RT), for T in K, R in cal/mol/K, t in seconds, and concentrations in (mol/cm3). • At 1.0 atm and 2000 K, the molar density of any perfect gas is 6.1  10-6 mol/cm3, so that

47. and • Therefore,

48. This is a simplified version of the Zelodvich simplification of the kinetics of the thermal NO formation reaction. • Using this simple relation, one may readily make up plot like Fig. 12.5 (next slide), which show the expected time-temperature relation for one specific starting gas composition.

50. More complex versions of the Zeldovich mechanism add another equation to the reaction list, N + OH ↔ NO + Hwith a resulting increase in mathematical complexity. • The result cannot be reduced to any simple plot like Fig. 12.5.