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Control of Nitrogen Oxides

朱信Hsin ChuProfessor

Dept. of Environmental EngineeringNational Cheng Kung University

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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.

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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.

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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.

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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.

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(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.

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We see that vehicles contribute almost half of the total.

Natural gas appears in this table, but not in the SO

2 table. The largest contributors to the “All other sources”

are nontransportation internal combustion engines and residential combustion.

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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.

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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.

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(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 i

s not a suitable fate for the nitrogen oxides.

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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.

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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.

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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.

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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

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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.

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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.

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From that condition it follows that

where ∆Go = standard Gibbs free energy changeK = equilibrium constantR = the universal constantT = 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).

(4)oG

ln KRT

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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.

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Example1 Calculate the equilibrium concentrations of NO a

nd 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).

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Solution:The definitions of the two equilibrium constants are:

Solving Eq. (3) for the equilibrium concentration of NO, we find

2

32 2

2 1 1/ 2

2

[ ] (5)

[ ][ ]

[ ]= (6)

[ ][ ]

NOK

O N

NOK

NO O

1/ 2

3 2 2[ ] [ ][ ]NO K N O

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Substituting values, including the value of K3 from Table 12.2, into this equation, we see

Solving Eq. (6) for [NO2], we find

1/ 2[ ] 0.0004[0.78][0.21] 0.0081 8100 NO ppm

1/ 22 1 2[ ] [ ][ ]NO K NO O

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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.

1/ 2 52[ ] 0.0035[0.0081][0.21] 1.3 10 13 #NO ppm

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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).

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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.

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(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.

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(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.

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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.

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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.

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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.

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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.

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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.

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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 ↔ 2NO2 ↔ 2OH2O ↔ H + OH

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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

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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 + N

and that N radicals can form NO by N + O2 ↔ NO + O

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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.

21/ 2

2 2 1/ 22

[ ] [ ][ ][ ] (7)

[ ]f b

d NO NOk N O k

dt O

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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).

3f

b

kK

k

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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

2 2 1/ 22

[ ] (8)

[ ] -[ ] [ ]b

e

kd NOdt

NO NO O

[ ] 1 ( ) (9)

[ ] 1 ( )e

NO exp t

NO exp t

1/22

2[ ] (10)

[O ]e bNO k

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Example 2 Estimate the concentration of NO in a sample con

taining 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.

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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

6 83

6 72 3

[ ] 6.1 10 0.003530 2.15 10

[ ] 6.1 10 0.04 2.44 10

e

molNO

cmmol

Ocm

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and

Therefore,

1 213

3

-8

-7 1/ 2

-91600 1 4.1 10 exp 4003

1.987 2000

2 2.15 10 4003 0.349

(2.44 10 )

b

molk

cm s

s

[ ] 1 exp[( 0.349 / )(1 )]0.173

[ ] 1 exp[( 0.349 / )(1 )e

NO s s

NO s s

[ ] 3530 0.173 610 #NO ppm ppm

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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.

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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.

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The Zeldovich mechanism shown here makes reasonably accurate predictions of the rate of formation of NO from N2 and O2 in the highest temperatures observed in flames, and in other high-temperature situations.

It predicts much lower NO concentrations than those observed in low-temperature flames.

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Before we discuss the prompt and fuel NO mechanisms that predominate at low temperatures, we explore the temperature and times in flames, to help understand what kinds of flames produce significant thermal NO and what kinds do not.

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2.2 Heating and Cooling Times

How much thermal NO is formed in a flame as well as how much is then converted back to N2 and O2 as the gases cool is a strong function of how fast the gases heat and cool in the flames.

Fig. 12.6 (next slide) shows some possible heating and cooling patterns for flames.

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The square wave of Fig. 12.6a leads to the easiest calculations, as shown in Example 2.

That example implicitly assumed that cooling after the flame is instantaneous, so that none of the NO formed in the flame is converted back to N2 and O2.

Nothing in this world is truly instantaneous!

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Example 3 Estimate the heating and cooling rates in the flame of an

ordinary gas stove. The peak flame temperature is about 2400oF and the

combustion time about 0.005 s. Solution:

If we assume that the temperature-time pattern is the triangle wave shown in Fig. 12.6b, we can say that the heating and cooling rates should be about the same, and

5 52400 689.3 10 5.2 10 #

0.0025

o o o odT T F F F C

dt t s s s

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The assumption of a symmetrical triangle temperature-time pattern is closer to reality than the square wave and still leads to relatively simple calculations.

If we take the viewpoint of a person riding with a parcel of gas through the flame (The Lagrangian view), we can say that:

where r is the rate of combustion and UA the product of the overall heat transfer coefficient and the applicable surface area.

heat added by combustion- heat lost to surroundings

=r ( ) (11)

P

combustion flame surroundings

dTV C

dth UA T T

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For the rising part of T-t pattern on Fig. 12.6c, the observed behavior is that the rate starts slowly, increases as the the temperature rises, and then slows to a stop as the last of the fuel is consumed.

After the peak is reached, the decline must be steeper than the rise was because the rise was proportional to the difference between the two terms on the right of Eq. (11) but the decline is proportional to the cooling term alone.

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Example 4 Repeat Example 2 for the strong simplifying

assumptions that (1) the gas heats from 293 to 2000 K in 0.5 s with a linear temperature increase, then cools from 2000 to 293 K in 0.5 s with a linear temperature decrease (Fig. 12.6b) and (2) the initial NO concentration is zero.

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Solution:Here we use the simple zeldovich mechanism and Eq. (8), which we know to be only approximate.

1707293 293 3414 for 0 < t < 0.5 s

0.5and

17072000 ( 0.5 )

0.5

2000 3414 ( 0.5 ) for 0.5 s < t < 1 s

K KT K t K t

s s

KT K t s

sK

K t ss

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If we consider the time interval 0.0 ≤ t ≤ 0.01 s, we see that the initial temperature is 293 K and the final temperature is 293 + 34 = 327 K.

For this time interval the average temperature is 310 K.

Using that value, we may estimate the equilibrium constant as

304340021.9 exp 5.5 10

1.987 310PK

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For simplicity we will proceed as if the oxygen content were the 4% in that example, independent of time.

Substituting this value of Kp in Eq. (5) and solving, we find

At this temperature and one atmosphere pressure, the molar density is (4.46 10-5) (273 / 310) = 3.93 10-5 mol/cm3, so that in concentration units [NO]e = 4 10-16 (3.93 10-5) = 1.6 10-20 mol/cm3.

30 0.5 16[ ] (5.5 10 [0.78][0.04]) 4 10eNO

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At this temperature, using the equation from Example 2, we have

13 51916004.1 10 exp 1.1 10

1.987 310bk

2 2 51 20 2 2

1/ 2 5 1/ 22

388

([ ] [ ] )[ ] 1.1 10 ([1.6 10 ] [0] )

[ ] [0.04 3.93 10 ]

/ 2.2 10

b ek NO NOd NO

dt O

mol cm

s

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Multiplying this by ∆t = 0.01 s, we conclude that during the first 0.01 s the concentration of NO went from zero to 2 10-90 mol/cm3.

This value is negligible, but we now proceed for the remaining 99 time steps, each of the same size, using a spread-sheet (e.g., Excel), and find the results shown in Fig. 12.7 (next slide). The final calculated NO concentration is 20 ppm. #

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From Fig. 12.7, we see that the calculated equilibrium concentration goes from 5.5 10-24 ppm (≈ 0) to the 3530 ppm in Table 12.3, as the temperature rises, and then back to zero as the temperature falls.

The reaction rate is practically zero, and hence the amount of NO actually formed is practically zero until about 0.44 s (T ≈ 1780 K), after which the rate of NO formation becomes significant.

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By the time the Equilibrium and calculated curves cross, the temperature is low enough that the reaction rate is practically zero, so that the NO concentration remains practically constant during the further cooling.

One speaks of the concentration being “frozen” at its higher temperature value.

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3. Prompt NO During the first part of combustion, the carbon-

bearing radicals from the fuel react with nitrogen by:

CH + N2 ↔ HCN + Nand several similar reactions involving the CH2 and C radicals.

The N thus produced attacks O2 to increase the amount of NO formed; and the HCN partly reacts with O2, producing NO, partly with NO, producing N2.

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There does not appear to be any simple theoretical treatment like the Zeldovich mechanism.

But it appears certain that the amounts of NO found in low-temperature flames, which are much larger than one would predict from the Zeldovich thermal mechanism, are actually mostly prompt NO.

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4. Fuel NO

Some coals as well as the hydrocarbons derived from oil shale can contain up to 2% N.

Most of the fuel nitrogen is converted in the flame to HCN, which then converts to NH or NH2.

The NH and NH2 can react with oxygen to produce NO +H2O, or they can react with NO to produce N2 and H2O.

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Thus, the fraction of the fuel nitrogen that leaves the flame as NO is dependent on the NO/O2 ratio in the flame zone.

Keeping the oxygen content of the gases in the high temperature part of the flame low significantly lowers the fraction of the fuel nitrogen converted to NO.

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The fraction of the fuel nitrogen that appears as NOX in the exhaust gas is estimated to be typically 20% to 50%, depending on furnace conditions and, to some extent, the chemical nature of the N in the fuel.

Fig. 12.8 (next slide) shows one schematic view of the interactions of the there mechanisms of NO formation.

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5. Noncombustion Sources of Nitrogen Oxides

The production and utilization of nitric acid lead to emissions of NO and NO2, as do some other industrial and agricultural processes, e.g., silage production.

Table 12.1 shows that such sources account for about 3% of NOX emissions in the U.S.

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6. Control of Nitrogen Oxide Emissions

There are two possible approaches to controlling NOX in combustion gases:(1) Modify the combustion processes to prevent the formation of NOX.(2) Treat the combustion gas chemically, after the flame, to convert the NOX to N2.

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For smaller industrial sources, like nitric acid plants, other control techniques are used, e.g., scrubbing with solutions of NaOH and KMnO4.

The strongly oxidizing KMnO4 quickly converts NO to NO2, allowing it to be captured by the alkaline NaOH.

This process requires an expensive electrochemical regeneration of the KMnO4; it is not suitable for large-scale combustion sources.

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6.1 Nitrogen Oxide Control by Combustion Modification

The formation of thermal NO is increased by increasing any of the following three variables: peak temperature, time at high temperature, oxygen content at the high temperature.

Two-stage combustion or reburning as shown in Fig. 12.6d is a good way for combustion modification.

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In the first stage, the maximum temperature is lowered because not all the fuel is burned, and the maximum temperature is reached when all the oxygen has been used up, so that there is not enough oxygen to form NO.

In the second stage, enough of the heat released in the first stage has been removed that the maximum temperature reached-in the presence of excess oxygen-is low enough that the formation of NO is small.

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The advantage of this approach is that it is cheap. The disadvantages are that it requires a larger fire

box for the same combustion rate and that it is difficult to get complete burning of the fuel in the second stage.

Fig. 12.9 (next slide) shows an alternative approach to designing a low-NOX burner.

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In normal operation the incoming gas is thoroughly mixed with air and recycled combustion products (Flue Gas Recirculation, FGR), with about 15% excess air.

This thorough mixing, together with excess air, prevents there ever being any fuel-rich part of the flame, in which prompt NO can form.

Fig. 12.10 (next slide) show some test results for this burner, burning natural gas with 15% excess air and varing degrees of air preheat and FGR.

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The dilution with recycled flue gas simply lowers the temperature at every part of the flame.

As the figure shows, it is possible to reduce the NOX to 10 ppm if one is willing to bear the extra cost of pumping 20 to 40% of the flue gas back into the burner.

However, compared to other methods of achieving 10 ppm–i.e., catalytic treatment of the flue gas-the cost of FGR is very reasonable.

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This burner is a moderately low-NOX burner when it is burning fuel oil.

Ultra-low NOX levels are not achieved since most fuel oils contain bound nitrogen that is oxidized in the flames to NOX.

Also, because there will be a fuel-rich region at the surface of evaporating oil droplets, there will be prompt NO formation.

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6.2 Nitrogen Oxide Control by Postflame Treatment

Most post flame treatment processes add a reducing agent to the combustion gas stream to take oxygen away from NO.

Almost any gaseous reducing agent can be used, e.g., CO, CH4, other HCs, NH3, various derivatives of NH3.

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In modern auto engines the reaction is

For power plants and other large furnaces a variety of reducing agents are used, of which the most popular is ammonia.

The desired reaction is:6 NO + 4 NH3 → 5 N2 + 6 H2O (12)

2 22 2 2platinum rhodium catalystNO CO N CO

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However, there is always some oxygen present, which can lead to reactions like: 4 NO + 4 NH3 + O2 → 4 N2 + 6 H2O (13)is which one ammonia molecule is required for each molecule of NO.

The NO2 is reduced by: 2 NO2 + 4 NH3 + O2 → 3 N2 + 6 H2O (14)

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Reactions (12) – (14) can be carried out over a variety of catalysts, of which zeolites and supported TiO2 are the most popular, at about 700oF (370oC), or simply in the gas stream in the part of the furnace where the temperature is between 1600o and 1800oF (870o – 980oC).

Below 1600oF the reaction rate is too slow. Above 1800oF the dominant reaction becomes

NH3 + O2 → NO + 3/2 H2Owhich increases rather than decreases the NO content of the gas.

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The catalytic processes are called selective Catalytic Reduction, or SCR.

The high-temperature ones, without catalysts, are called Selective Noncatalytic Reduction, SNCR.

There are experimental processes under development that adsorb both SO2 and NO onto an activated carbon adsorbent.

They are then stripped off one at a time and treated, making S and N2 the final products.