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Ind. Eng. Chem. Res. 1992, 31, 777-780 actor. J. Appl. Polym. Sci. 1981,26, 3179-3204. Ravindranath, K.; Mashelkar, R. A. Modeling of Poly(ethy1ene TereDhthalate) Reactors: 11. A Continuous Transesterification Process. J. Appl. Polym. Sci. 1982,27, 471-487. Tomita, K.; Ida, H. Studies on the Formation of Poly(ethy1ene Terephthalate): 2. Rate of Transesterification of Dimethyl Terephthalate with Ethylene Glycol. Polymer 1973, 14, 55-60.
Zaborsky, L. M., 11. Determination of Polyester Prepolymer Oligomers by High Performance Liquid Chromatograph. Anal. Chem. 1977,49 (8),1166-1168.
Received for review September 6 , 1991 Revised manuscript received November 5, 1991 Accepted December 3,1991
Removal of NO, from Flue Gases Using the Urea Acidic Process: Kinetics of the Chemical Reaction of Nitrous Acid with Urea Alain Lasalle,t Christine Roizard,*vtNoel Midoux,t Pierre Bourret,: and Pierre J. D y e d Laboratoire des Sciences du G&nie Chimique and Laboratoire de Physico-Chimie Zndustrielle, CNRS-ENSIC-INPL, BP 451 54001 Nancy Cedex, France, and Socrematic S. A., BP 756 95004 Cergy-Pontoise, France
This study deals with the removal of nitrogen oxides from flue gases using the acidic urea process. The chemical hydrolysis of nitrous acid, which leads to NO formation, is avoided by nitrous acid reaction with urea. Products of this reaction are gases, e.g. C02and N2, which can then be directly released into the atmosphere. The aim here is to determine the kinetic parameters of the chemical reaction of nitrous acid with urea. Experiments are performed in a closed stirred reactor. The manometric method (measurement of the pressure versus time curve) leads to the determination of the concentration of H N 0 2 and then to the chemical rate versus time. Operating parameters are the concentration of urea (333-3330 mol m-3), the pH (0.75-1.25), and the temperature (3-40 "C). The experimental results are as follows: the order relative to nitrous acid is 1;the rate constant decreases with pH; the influence of temperature on the rate constant can be expressed by (pH = 1) k = 1.82 X lo8 exp(-(60400/RT)) (SIunits).
Introduction Nitrogen oxides, NO,, are among the major pollutants in the environment and thus have to be removed from industrial waste gases. Several dry or wet processes have been proposed previously. Joshi et al. (1985) and Jethani et al. (1990) give critical reviews of the wet processes. Among these, the aqueous acidic urea process is particularly interesting because urea is a cheap reactive and because products of reactions in the liquid phase are gases, i.e. carbon dioxide and nitrogen, which can be directly released into the atmosphere. This process has been patented by Warshaw (1971) in the USA and by Dyens (1985) in France. Absorption of NO, gases is a very complex system due to the numerous chemical reactions which can occur. NO, gas is a mixture of several components, NO, NO2,N204, and N203,and the following reactions occur in the gas phase: 2N0 + 02 2N0z (1)
-
-
2N02 N204 NO + NOz N203 NO + NO2 + HzO
Q
2HN02
(2) (3) (4)
The different gaseous components are absorbed into the liquid phase, with the exception of NO due to ita low solubility. In the case of pure water or aqueous solutions + Laboratoire des Sciences du GBnie Chimique, CNRS-ENSIC-INPL. Laboratoire de Phvsico-Chimie Industriele, CNRS-ENSICINPL. 1 SOCREMATIC S. A.
*
--
of nitric acid, the following chemical reactions occur in the liquid phase: 2NOZ+ HzO HN02 + HN03 (5) Nz04+ HzO HN02 + HNO, (6) N203+ HzO 2HNOZ (7) 3HNOZ HNO, H20 + 2NO (8) Reaction 8 shows that the decomposition of nitrous acid leads to the formation of NO, which desorbs into the gaseous phase. In the case of the acidic urea process, nitrous acid reacts with urea by the following mechanism: 2HN02 + NH2CONH2 2N2 + C02 + 3H20 (9)
-
Q
+
No fundamental study concerning this acidic urea process has been reported so far, as pointed out by Jethani et al. (1990). The aim here is to determine the kinetic parameters of reaction 9. Since the chemistry is very complex, only a limited range of parameters will be considered; however the considered range includes values of interest in future process design. The pH varies from 0.75 to 1.25 and the concentration of urea from 333 to 3330 mol m-,. Moreover the apparent kinetics of this reaction is assumed to be given by a relation of the form r = k[HN02]m[~rea]n where m,n are the apparent orders relative to nitrous acid and urea, respectively; k is the rate constant; and r is the kinetic rate. 1. Chemical Considerations The reaction of nitrous acid with urea occurs when the liquid phase is acidified with a strong acid (pH < 2) and when urea is in excess in order to avoid reaction 8 (Grig1992 American Chemical Society
778 Ind. Eng. Chem. Res., Vol. 31, No. 3, 1992
Inert gas (N inlet
Figure 2. Total pressure P ( E a ) vs time t ( 5 ) experimental curve: pH = 1; T = 15.5 "C; [urea] = 333 mol m-3; [NaN0210= 5.87 mol m-3.
baffles
Figure 1. Schematic diagram of experimental setup.
nard et al., 1939). In fact, the chemical scheme has been reported by Werner (1917) as "02 + NHzCONHz N2 + HNCO + 2H20 (10) "02
--
+ HNCO
N2 + C02 + H2O
(11)
However, no chemical data dealing with kinetics or chemical equilibrium of these reactions has been reported. Considering the pH range of this study, HNCO can only react with nitrous acid; it cannot be hydrolyzed (Werner, 1917). Thus it seems reasonable to consider only the total reaction 9. In order to avoid the following possible ionic dissociation reactions "02 * H+ + NO2(12) NH03
-
H+ + NO3-
2. Experimental Method For measurement of the kinetics of reaction 9, the manometric method was used; the total pressure versus time in a closed agitated reactor was measured. The pressure is related to the products of the reaction, in this case C02 and NP. Due to the complexity of the prmss, nitrous acid is introduced directly in the form of sodium nitrites, which dissociate completely in aqueous solutions:
-
Na+ + NOz-
2A
mainly of a closed stirred reactor. The stirring system consists of two turbines, one in the gas and one in the liquid phases. The rotational stirring speed was fixed at a high value (1300 rpm) in order to avoid mass-transfer resistance (Mercadier, 1989). The temperature is kept
+U
2B
2N2 + C02 + 3H20 (9) C + 3H20
several assumptions have to be made. (a) No side reactions occur (cf. section 1). (b) Carbon dioxide is in thermodynamic equilibrium, and Henry's law is assumed (coefficient H). (c) Nitrogen, N2, is not soluble in the urea solution. (d) The liquid and gas volumes are constant; gases are considered as ideal. Since the gas phase is assumed ideal, the total number of moles of gas produced per unit time is - = - - =vg - dp dng
dt
RT dt
dnBg
dt
+ -dnCg dt
(15)
The stoichiometry of reaction 9 leads to
and since nitrogen is assumed not soluble in the urea solution dnBg dt
dnB dt
-=-=--.
dnA dt
(17)
The produced C02 desorbs partially and is partially dissolved into the liquid phase. Using Henry's law and eqs 15-17, one obtains
(14)
2.1. Experimental Setup. A schematic diagram of the experimental setup is shown in Figure 1. It consists
- +
2HN02 + NH2CONH2
(13)
the experimental pH values were fixed between 0.5 and 1.3 since the pK, of nitric acid is -1.4 and the pK, of nitrous acid is 3.3 at 25 "C (Charlot, 1983). Obviously, it is clear that outaide this pH range, the kinetics follow a more complex chemical scheme. In addition to the above reactions, the hydrolysis of urea should be added.
NaN02
constant in the reactor by a thermostatic bath, controlled by a thermometer. A pH probe and meter allow the measurement of this parameter. Finally a pressure sensor measures the total pressure in the gas phase continuously. An aqueous urea solution, acidified with nitric acid, is placed initially into the reactor; then a stream of inert nitrogen ensures the stripping of oxygen and carbon dioxide off the liquid. A low quantity of sodium nitrite is then added into the reactor at t = 0; the gas phase consists then of pure nitrogen with Po= 106 Pa (absolute pressure) = 0 (relative pressure). Figure 2 is an example of the pressure history in the gas phase. 2.2. Determination of the Rate of Reaction. For determination of the kinetics of the reaction
dt
(18)
Finally, with the addition of eqs 18 and 17, the rate of reaction 9 as a function of the total pressure is obtained:
Ind.Eng.Chem. Res., Vol. 31, No. 3, 1992 779 -3.5 I
I
I
I
I
-4
I
-4.5
-5 -5.5 -6
4-
-6.5
2 -
.'I -1.5
0. 0 1 0 0 u ) o 3 0 0 4 0 0 5 0 0 6 0 0 7 0 0 8 0 0
Figure 3. Nitrites concentration [HN02] (mol rn-9 va time t (a) curve: pH = 1; T = 15.5 O C ; [urea] = 333 mol m-3; [NaN0210= 5.87 mol m+.
-1
0
-0.5
0.5
1.5
1
2
2.5
Figure 5. Determination of the order relative to nitrous acid m, In r versus In [HN02]curve: pH = 1; T = 15.5 "C;[urea] = 333 mol m-3; [NaN0210= 5.87 mol m-3. 0.03 >
0.03
0.025
0.025
0.02
0.02
0.015 0.01
0.015
0.005
0.01 0 0
0.005 0
0
100
200
300
400
600
500
700
Figure 4. Reaction rate r (mol m-3 8-l) va time t (a) curve: pH = 1; T = 15.5 O C ; [urea] = 333 mol m-3; [NaN0210= 5.87 mol m-3.
where a is a constant due to the above assumptions and is defined as a=
(4 + -)/(5+ 5) VIRT V P
RT
(
4
6
8
10
12
Table I. Experimental Results of the Order Relative to Nitrous Acid, m ,and of the Rate Constant. k ["02lo, T,O C pH mol m-3 [ureal, mol m-3 m k. a-l 20 1.05 10.67 416.67 1.06 0.00402 833.33 20 1.05 10.71 0.00391 1.09 1250.00 20 1.05 10.63 1.17 0.003 76 20 1.05 10.71 0.00418 1666.67 1.00 2500.00 10.91 20 1.05 1.15 0.00378 20 1.05 10.85 0.00379 3333.33 1.06 11.74 21.2 1.05 333.33 1.15 0.00387 20.5 1.1 10.00 1333.33 1.04 0.00415 2000.00 1.04 11.86 20.5 1.00 0.00418 1666.67 22.4 1.09 10.10 1.14 0.004 52 20.7 1.03 11.74 1.06 0.00425 3000.00 ~
(20)
H
Equation 19 can be integrated from the initial condition nA= no and P = Poto the measured final values at tf, nA = 0 and P = Pf;this leads to the expression of nA n A = n o l--
2
Figure 6. Determination of the rate constant k, r (mol m-3 8-l) va [HN02] (mol m-3) curve: pH = 1;T = 15.5 "C; [urea] = 333 mol m-3; [NaNO& = 5.87 mol m-3.
;2))
The a value is then obtained: a = -
Pf - Po Equation 19 can be rewritten as
For reaction 9 to occur, the liquid phase must be acidic; nitric acid was used to ensure this. Experiments were performed using hydrochloride acid to verify that nitric acid does not influence the reaction mechanism; the same results were obtained.
3. Kinetics Determination at pH = 1 and T = 20
"C Using the experimental pressure versus time curve, it is then possible to calculate the rate r using eq 23 and the nitrite concentration from eq 21. Examples of such curves are presented in Figures 3 and 4. 2.3. Preliminary Experiments. Total mass balances, using the stoichiometry of reaction 9 and assuming all the nitrites have reacted, allow us to determine the produced gas mole number and the final total pressure in the gas phase. The obtained values are in good agreement with the experimental data. Moreover, some samples of the gas phase were analyzed by chromatography. It results that carbon dioxide and nitrogen partial pressures were in the ratio of 1-2, obviously after subtraction of the initial pressure of nitrogen, which is in accordance with reaction 9. One can thus conclude that reaction 9 is clearly the only one occurring in such a system when the pH is between 0.75 and 1.25.
As previously stated, the apparent kinetics is assumed to follow the Arrhenius law: r = k[HNO,]"[urea]" The concentration of urea must be in excess in order to avoid reaction 8; this concentration is assumed to be constant during the experiment. Hence a plot of In r vs In [HNO,] gives a straight line of slope m,the order relative to HN02, and the intercept with the y-axis In [k[urea]"] (Figure 5). For this series of experiments, the parameters were the initial concentration of nitrites (from 10 to 20 mol m-3) and the concentration of urea (from 330 to 3330 mol m-3), the pH was fixed at a value of 1,and the temperature was 20 "C. Owing to the large excess of urea, the order relative to urea is obviously zero. The experimental m values are shown in Table I. It is clear that m = 1. It is then possible to plot directly r versus [HN02], whose slope gives values of k (Figure 6). The
780 Ind. Eng. Chem. Res., Vol. 31, No. 3, 1992
Conclusions The kinetics parameters of the chemical reaction of nitrous acid with urea were determined experimentally in a closed stirred reactor. The experimental results are as follows: the order relative to nitrous acid is 1;the rate constant decreases with pH; the influence of temperature on the rate constant can be expressed by (pH = 1)
0.005
0.00 0.004
k
0.00 0.003 0.00
0.002
0.002.
.
.
.
.
.
( 6:!3 (SI units)
.
k = 1.82 X lo8 exp --
0.75 0.8 0.85 0 . 9 0.95 1 1.05 1.1 1.15 1 . 2 1.25
PH
Figure 7. Influence of pH on the rate constant k
(9-l).
-4 -4.5 -5
-5.5
Ln k
-6
-6.5 -7
It is clear that additional experiments should be performed to quantify the influence of pH on the rate constant. If the chemical reaction rates are known, it is then possible to study the acidic urea process and to measure and model the coupled mass-transfer/chemical reaction phenomena occurring in the gas and in the liquid phases. Nomenclature E = activation energy, J mol-' H = Henry's constant, Pa m3 mol-' k = rate constant, (mol m-3)1-m-ns-l m = order relative to nitrous acid n = order relative to urea nj = number of moles of species i , mol P = total pressure, Pa r = rate of reaction, mol m-3 s-l R = gas constant, J mol-' K-' t = time, s T = temperature, K V = volume, m-3 CY = constant defined by eq 6, mol Pa-'
5
- 7-*85 0.00035 0.00037 0.00039 0.00041 0.00043 0.00045
1 IRT Figure 8. Influence of temperature on the rate constant k (s-l). Experimental results of In k versus 1/RT (mol J-l).
obtained values are reported in Table I. The kinetic constant thus obtained at pH = 1 and T = 20 "C is
k = 4.0 X
s-l
4. Influence of pH
Several experiments were performed at different pH values in the range of 0.75-1.25. Results of the rate constant k versus pH comparison are shown in Figure 7. It is clear that the rate constant decreases with pH. Note that pH values are constant during one experiment; the pH is fixed by the initial concentration of nitric acid. Moreover, some experiments conducted at higher pH values (1.5 and 2) show clearly that the additional chemical equilibrium reactions (12) have to be taken into account, the order relative to nitrous acid is no longer 1, and pH varies with time during the experiments. The experimental technique has to be modified if the kinetics parameters at such pH values are to be determined. 5. Influence of Temperature
A temperature range of 3-40 "C was investigated; the pH value was fixed at 1. Considering the Arrhenius law
k = k, exp( -
&)
a plot of In k vs 1/RT gives a straight line of slope E, the activation energy, and the intercept with the y-axis is k, (Figure 8). The experimental values obtained are
E = 60400 J mol-'
and k, = 1.82 X lo8 s-'
Subscripts A = nitrous acid B = nitrogen C = carbon dioxide f = final time g = gas 1 = liquid U = urea 0 = initial time
Literature Cited Charlot, G. Les rlactions chimiques en solution aqueuse et caract6risation des ions; Masson: Paris, 1983; pp 354-362. Dyens, P. J., Socrematic S. A. Proc6dB et dispositif #elimination par voie humide des oxydes d'azote d'un effluent gazeux. French Patent 85 13 146, 1985. Grignard, V.;Dupont, G.; Locquin, R. Trait6 de chimie organique; Masson: Paris, 1939; Vol. 14, p 39. Jethani, K. R.; Suchak, N. J.; Joshi, J. B. Selection of reactive solvent for pollution abatement of NOx. Gas Sep. Purif. 1990, 4 , 9. Joshi, J. B.;Mahajani, V. V.; Juvekar, V. A. Absorption of NOx gases. Chem. Eng. Commun. 1985,33, 1. Mercadier, J. Cinetique de la reaction des nitrites avec l'urBe en vue de l'Bpuration d'effluents gazeux contenant des oxydes d'azote. DEA Dissertation, Institut National Polytechnique de Lorraine at Nancy, France, 1989. Warshaw, A., Chemical Construction Corp. Removal of nitrogen oxides from a gas stream. US Patent 3 565 575, 1971. Werner, E. A. The constitution of carbamides. Part IV. The mechanism of the interaction of urea and nitrous acid. J. Chem. SOC.London 1917,112, 863.
Receiued for review July 22, 1991 Revised manuscript received November 13, 1991 Accepted December 3, 1991
