Kinetics of Cyanate Decomposition in Alkaline Solutions of High Ionic

Ind. Eng. Chem. Res. 2004, 43, 4815-4821

4815

APPLIED CHEMISTRY Kinetics of Cyanate Decomposition in Alkaline Solutions of High Ionic Strength: The Catalytic Effect of Bicarbonate Nikolai DeMartini,* Dmitry Yu. Murzin, Mikael Forsse´ n, and Mikko Hupa Process Chemistry Group, A° bo Akademi University, Piispankatu 8, Turku/A° bo, Finland 20500

This paper clarifies the catalytic effect of bicarbonate on the decomposition of cyanate in highly alkaline solutions of high ionic strength at 94 °C containing sodium cyanate (0.019 M), sodium hydroxide (0.1-0.9 M), and sodium carbonate (0-1.1 M). Activity coefficients were calculated to account for nonideality in determining the concentration of bicarbonate in solutions with a pH above pKa,HCO3- . A rate equation is presented, and parameter estimation was used to obtain values for the rate constants based on the experimental data. Finally our data and literature data was used to derive the temperature dependence of the relevant rate constants and compare the predictions of the kinetic model against the experimental results with three kraft pulp mill green liquors. Introduction Ammonia is formed from cyanate in the strongly alkaline (pH > 12) green and white liquors in the chemical recovery cycle of kraft pulp mills and can lead to NH3 and NOx emissions depending on how the vent gases and condensates are handled.1-5 These solutions have both a high ionic strength (>6) and high carbonate concentrations (1-1.6 M). Nitrogen containing compounds enter the pulping process with the wood and are extracted with the lignin during pulping.6 About 1015% of this organic nitrogen forms ammonia with the remaining 85-90% contained in various organic compounds.7 The organic nitrogen compounds exit with the black liquor, and approximately 20-30% of this nitrogen is converted into the inorganic nitrogen compound cyanate during the char burning stage of black liquor combustion.8 This cyanate then reacts in the green and white liquors during recausticizing to form ammonia. Cyanate is the anionic form of cyanic acid. There are two isomers of cyanic acid, viz. (H-O-C≡N) and isocyanic acid (HNdCdO). In aqueous solutions both isomers exist, but isocyanic acid is the predominant species.9 Cyanate is thus likely a resonance hybrid of -NdCdO and N≡C-O-. In the text the acid will be referred to as isocyanic acid and the anion as cyanate. The notation used for cyanate will be OCN-. Previous studies focused on the formation of NH3 from cyanate as a step in the decomposition of urea in aqueous solutions covering the pH range from 1 to over 13 and included the catalytic effects of several anions10-14 including bicarbonate. Bicarbonate is found in low concentrations in the chemical recovery streams following acid-base equilibrium. These works have * To whom correspondence should be addressed. Fax: +358 2 2154962. E-mail: [email protected].

agreed on the form of the kinetic equation expressing the decomposition of isocyanic acid plus cyanate in aqueous solutions of various pH. It can be written as eq 1.

(

)

[H3O+]S d[S] d[NH3] + ) ) k1[H3O ] + dt dt KOCNH + [H3O+]

(

k2

[H3O+]S +

KOCNH + [H3O ]

) ( + k3

KOCNH × S

KOCNH + [H3O+]

)

(1)

The total substrate, S, is the sum of the molar concentration of isocyanic acid and cyanate, while the terms ([H3O+]/KOCNH + [H3O+]) and (KOCNH/KOCNH+ [H3O+]) correspond to the portion of substrate that is isocyanic and cyanate, respectively. The constant KOCNH is the acid dissociation constant for HNCO (KOCNH ) 2.2 × 10-4 at 20 °C).15 The rate constants k1 is for the hydronium ion catalyzed hydrolysis of isocyanic acid, and k2 is the rate constant for the direct hydrolysis of isocyanic acid. These two reactions dominate the kinetics in acidic (pH 0-2) and neutral conditions (pH 3-8), respectively. The third rate constant, k3, represents the direct hydrolysis of cyanate. Vogels et al. offer an alternative reaction: the hydroxide catalyzed hydrolysis of isocyanic acid.13 It is not clear which of these two reactions is responsible for ammonia formation in alkaline solutions, but for the purposes of a kinetic description it is possible to describe the kinetics with the assumption that it is cyanate that is hydrolyzed in alkaline conditions. The third term dominates the rate of reaction at high pH (>11). However, differing conclusions were drawn in the literature as to the catalytically active species of the carbonic acid system in highly alkaline solutions, and different kinetic expressions were offered for this reaction.10,13,14 A possible mechanism for the catalytic effect

10.1021/ie034189+ CCC: $27.50 © 2004 American Chemical Society Published on Web 06/30/2004

4816 Ind. Eng. Chem. Res., Vol. 43, No. 16, 2004 Table 1. Initial Concentrations of Anions in the Experimental Solutions

Table 2. Analyzed Concentration of Major Ions in the Mill Green Liquors

experiment

OH(mol/L)

CO32(mol/L)

OCN(mol/L)

pHa

mill

carbonate (mol/L)

hydroxide (mol/L)

sulfur (mol/L)

chloride (mol/L)

sodium (mol/L)

1 2 3 4 5 6 7 9 10 11 12 13 17

0.1 0.1 0.1 0.2 0.9 0.1 0.2 0.9 0.9 0.5 0.5 0.1 0.5

0 0.07 0.15 0.07 0.07 1.1 1.1 0.5 1.1 0.15 0.5 0.5 0

0.019 0.019 0.019 0.019 0.019 0.019 0.019 0.019 0.019 0.019 0.019 0.019 0.019

11.2 11.2 11.1 11.5 12.1 11.1 11.4 12.1 12.0 11.8 11.8 11.1 11.9

A B C

1.28 1.58 1.29

0.32 0.31 0.11

0.77 0.51 0.83

0.041 0.097 0.019

3.19 4.48 4.56

a Calculated based on the thermodynamically determined activity of H+.

of bicarbonate is the formation of a carbamoyl monoanion intermediate which rapidly decomposes to carbamate and CO2 with the carbamate further decomposing to NH3 and CO2. Due to the rapid decomposition of the carbamoyl intermediate, it cannot be analyzed directly.13 This carbomoyl intermediate would be formed by bonding of one of the oxygen atoms in the bicarbonate anion to the carbon atom in cyanate or isocyanic acid molecule similar to the formation of carbamic acid by hydrolysis of isocyanic acid which occurs in the pH range of 2-8.11,13 This paper gives a clarification of the catalytically active species as well as a kinetic expression for the catalyzed hydrolysis in alkaline solutions, which takes into account both relevant earlier data and the experimental data presented in this work. Experiments were carried out at 94 °C with aqueous solutions of cyanate (0.019 M), hydroxide (0.1-0.9 M), and carbonate (0-1.1 M) to clarify the effect of these anions on the reaction rate of cyanate decomposition. These conditions were chosen both because they extended the concentrations of these ions beyond the previous research and because they were directly applicable to the kraft pulp mill recovery solutions. Nonideality was accounted for in the catalyzed reaction kinetics by calculation of activity coefficients. The ratio of relevant activity coefficients was also plotted as a function of the ionic strength, to provide a method for approximating the activity coefficient ratio without detailed chemical equilibrium calculations. Parameter estimation was performed to obtain rate constants for the uncatalyzed and catalyzed reactions in highly alkaline solutions. This was done for both our data and a part of the literature data for obtaining the temperature dependence of the relevant rate constants. Finally, the predictive power of the kinetic model was tested against new experimental data of three kraft pulp mill green liquors at 95 °C.

through an opened sample port, which took approximately 90 s. The sample was then transferred to a plastic bottle that was sealed and cooled in an ice bath for approximately 5 min, until it reached room temperature. In the second method, the sample port was replaced with a plug containing a sealed sample line and valve. Before the sample was pulled, the line was purged with air, and then the sample was drawn. The sample bottle was then handled as before. Any effect of the cooling time on determining the rate constants is negligible due to the slow rate of the reaction. At room temperature, the sample was considered stable, and samples were analyzed within 48 h. Green liquors from the smelt dissolving tank were pulled at three kraft mills pulping: a softwood mixture (80% spruce/20% pine), eucalyptus, and a hardwood mixture (80% birch/20% aspen) at mills A, B, and C, respectively. The concentrations of the main ions in these green liquors are given in Table 2. Analytical. Cyanate was analyzed using ion chromatography with chemically suppressed conductivity detection following Metrohm application S2.2,16 Carbonate in the mill green liquors was determined according to the Scandinavian method SCAN-N 32:98 and OHby SCAN-N 30:85. Total sulfur was determined according to Scandinavian method SCAN-N 5:83. Sodium was determined by ICP-AES. Results and Discussion Catalytic Effect of Bicarbonate. Experiments 1 and 17 which initially contained only cyanate and hydroxide indicated that there was no effect of increasing the hydroxide concentration above 0.1 M, Figure 1. However, at a hydroxide concentration of 0.1 M, the addition of sodium carbonate was found to increase the rate of cyanate decomposition (experiments 1, 2, 3, 6, and 13), Figure 2. As the pH is increased for the same initial sodium carbonate concentration, the rate of cyanate decomposition slows (experiments 6, 7, and 10), Figure 3. These results can be explained qualitatively with a logarithmic diagram of acid-base equilibrium for the carbonic acid system, Figure 4, which can be derived as described by Ha¨gg17 and Sille´n.18 Above the pKa for

Methods Experimental. The kinetic experiments were carried out in either a 1-dm3 or 3.7-dm3 Pyrex glass vessel with stir rod and thermometer. The reactor was heated in an oil bath. Sodium carbonate (Na2CO3, Baker), sodium cyanate (NaOCN, Aldrich), and aliquots of 0.5 or 2.0 N sodium hydroxide solution (NaOH) were added to ultrapure water to obtain the initial solution for each experiment (Table 1). Two sampling methods were used. During the first experiments, the sample was drawn with a glass pipet

Figure 1. Cyanate concentration vs time for [CO32-]i ) 0 M and [[OH-]i ) 0.1 M; ∆ [OH-]i ) 0.5 M.

Ind. Eng. Chem. Res., Vol. 43, No. 16, 2004 4817

Figure 2. Cyanate concentration vs time for [OH-]i ) 0.1 M; and [[CO32-]i ) 0 M; ×[CO32-]i ) 0.07 M; ∆[CO32-]i ) 0.15 M; O[CO32-]i ) 0.5 M; 0[CO32-]i ) 1.1 M. Figure 4. Logarithmic diagram for the acid base equilibrium of the carbonic acid system at 95 °C calculated as demonstrated by Ha¨gg17 and Sille´n;18 [H2CO3] + [HCO3-] + [CO32-] ) 1.1 M.

Figure 3. Cyanate concentration vs time for [CO32-]i ) 1.1 M and [[OH-]i ) 0.1 M; ∆[OH-]i ) 0.2 M; O[OH-]i ) 0.9 M.

bicarbonate, any increase in pH results in a decrease in the bicarbonate concentration. Similarly, any increase in the total carbonate concentration will result in an increase in the bicarbonate concentration for the same pH. Following Jensen14 it is reasonable to assume that eq 1 should contain an additional term accounting for the catalytic effect of bicarbonate on the reaction rate. As the experiments in the present study were conducted in highly alkaline conditions, eq 1 is then modified to

-

(

)

KOCNH × S d[OCN-] + ) k3 dt KOCNH + [H3O+] KOCNH × S [HCO3-] (2) kHCO3KOCNH + [H3O+]

(

)

The first two terms from eq 1 have been dropped as they are very small at the high alkalinity of the experiments in this study. The rate constants k3 and kHCO3- are for the uncatalyzed and bicarbonate catalyzed hydrolysis reactions, respectively. At a pH above pKa,HCO3- (10.33 at 20 °C),17 the bicarbonate concentration can be determined from acid-base equilibrium

[HCO3-] )

KH2O [CO32-] fCO32KHCO3- [OH-] f - f OH HCO3-

(3)

where Ki is the acid dissociation constant, [Xi] is the molar concentration, and and fi is the activity coefficient of species i. The acid dissociation constants for bicarbonate and water are given in Figure 5 for the temperature range 0-100 °C.

Figure 5. Acid dissociation constants, KHCO3- and KH2O from 0 to 100 °C.

Calculation of Activity Coefficients. The activity coefficients for the anions CO32-, HCO3-, and OH- in each experiment were calculated with the program ChemSheet, which uses Gibbs energy minimization calculations to determine the equilibrium concentrations as described by Eriksson and Hack.19 The calculations for the aqueous species employ the Pitzer model.20 The database used for these calculations was compiled at A° bo Akademi University and included aqueous ions and gaseous and solid compounds relevant to green liquor solutions. The parameters for the full Pitzer model were obtained from the thermodynamic database FACT 3.05.21 Calculation results are presented in Table 3. The calculated activity coefficients for the carbonate and bicarbonate anions decrease with increasing ionic strength and begin to level at an ionic strength of 3.5. This is consistent with the general trend of the mean activity coefficient for Na2CO3 as determined with vapor pressure measurements at 95 °C by Taylor.22 Similarly the calculated activity coefficient for the hydroxide anion generally decreases until an ionic strength of 2 is reached and then remains nearly constant. Harned23 found that for an aqueous solution of NaOH at 25 °C, the mean activity coefficient for NaOH decreased until an ionic strength of 0.8 and then increased again. The thermodynamic model was used to calculate the mean activity coefficients for the experimental matrixes in Taylor22 and Harned23 (Figure 6). Calculations for the mean activity coefficient of NaOH were also calculated at 95 °C, and they indicate that the increase in mean

4818 Ind. Eng. Chem. Res., Vol. 43, No. 16, 2004 Table 3. Activity Coefficients and Ionic Strength for the Initial Compositions of the Synthetic Liquors at 94 °C and Three Mill Green Liquors at 95 °Ca experiment

fCO32-

fOH-

fHCO3-

fCO32-/ (fOH-×fHCO3-)

ionic strength

1 2 3 4 5 6 7 9 10 11 12 13 17 mill A mill B mill C

0.26 0.14 0.094 0.12 0.051 0.017 0.017 0.025 0.014 0.060 0.030 0.037 0.096 0.010 0.0096 0.011

0.75 0.64 0.58 0.62 0.59 0.48 0.48 0.52 0.51 0.57 0.51 0.5 0.64 0.49 0.49 0.48

0.71 0.56 0.48 0.53 0.43 0.26 0.26 0.32 0.25 0.43 0.33 0.34 0.53 0.22 0.21 0.22

0.49 0.39 0.33 0.35 0.2 0.14 0.13 0.15 0.11 0.24 0.18 0.21 0.28 0.097 0.091 0.10

0.12 0.33 0.57 0.43 1.13 3.42 3.52 2.42 4.22 0.97 2.02 1.62 0.52 6.50 6.70 6.50

a

Calculated using ChemSheet.

activity coefficient is less pronounced at higher temperatures. For the mean activity coefficient of Na2CO3, the calculated values are 11% lower at an ionic strength of 0.3 and 36% at an ionic strength of 4.5. For NaOH, the discrepancy is much lower being 5% at an ionic strength of 3.1. These findings indicate that actual ionic activity coefficients could be higher than those calculated, though this is not conclusive as the experimental solutions tested here are concentrated mixtures of both NaOH and Na2CO3. Similar values were obtained for the ratio of activity coefficients in the three industrial green liquors as listed in Table 3. As the Pitzer model takes into account ionic interactions, the individual ionic activity coefficients may not be the same for solutions of the same ionic strength but different composition as seen for mill liquors A and C. The ionic strength was also determined from the molar concentration of the ions (Ci) and the respective charge (zi) using eq 4:

I)

1 2

∑Cizi2

(4)

The ratio of activity coefficients in eq 3 follows an almost

Figure 7. The ratio of activity coefficients from eq 6 plotted against ionic strength at 94 °C for the solutions and 95 °C for the green liquors.

exponential decay as a function of ionic strength for the solutions and green liquors studied in this work (Figure 7). Parameter Estimation. Parameter estimation was carried out on our data and Jensen’s high alkaline data with carbonate added at 94 °C,12 because even at the relatively low concentrations of carbonate (0.07 and 0.15 M), the activity coefficients for the relevant ions are much less than one and must be accounted for. For parameter estimation purposes primary concentration vs time data from all the experiments was used, and the following modified variant of eq 2 was tested.

-

[CO32-] d[OCN-] ) p1[OCN-] + p2p3[OCN-] dt [OH-]

(5)

The parameters p1, p2, and p3 represent k3 (rate constant in alkaline conditions), kHCO3- (rate constant for the bicarbonate catalyzed reaction), and the ratio of activity coefficients, fCO32-/fOH- fHCO3-, respectively. Because the amount of cyanate is small compared to the concentration of hydroxide and carbonate, the ratio of activity coefficients was found to vary by only 1.1% for experiments 1 and 2 and by less than 1% for more concentrated solutions. Thus, for the purpose of parameter estimation, this ratio was held constant, while the changes in the carbonate and hydroxide concentrations were calculated for each experimental data point. The change in bicarbonate concentration through an experi-

Figure 6. Mean activity coefficient (f++) for O NaOH, 25 °C data Harned;23 × NaOH, 25 °C, calculated using ChemSheet; b NaOH, 95 °C, calculated using ChemSheet; ] Na2CO3, 95 °C data Taylor;22 ∆ Na2CO3, 95 °C, calculated using ChemSheet.

Ind. Eng. Chem. Res., Vol. 43, No. 16, 2004 4819 Table 4. Parameter Values and Statistical Analysis for (A) Jensen’s Experimental Data (Runs 5-714 and Run 5012); (B) Our Experiments 1-4; (C) All of Our Experiments Listed in Table 1; and (D) All of the Experiments Listed in Table 1 but with p1 Fixed data set parameter estimation values (min-1)

p1 p2 (L mol-1 min-1) total SS (corrected for means) residual SS SE of estimate degree of explanation (%)

A

B 10-2

C 10-2

0.166 × 0.209 0.2046 × 10-3 0.5108 × 10-6 0.1459 × 10-3 99.75

0.170 × 0.241 0.3196 × 10-3 0.2642 × 10-5 0.3466 × 10-3 99.17

D 10-2

0.158 × 0.371 0.1321 × 10-2 0.2690 × 10-4 0.5350 × 10-3 97.96

0.170 × 10-2 0.33 0.1423 × 10-2 0.3901 × 10-4 0.6246 × 10-3 97.26

ment was thus accounted for, though this change was slight as the amount of carbonate formed from cyanate is small in comparison to the initial carbonate concentration. The differential equation was solved numerically with the backward difference method using the software Odessa. The differential equation solver operated under a parameter estimation routine that minimized the following objective function, the residual sum of squares

Q)

∑t (cA(t) - cˆ A(t))2

(6)

where cA and cˆ A denote the experimental and predicted concentrations at a given time, t, respectively. A hybrid simplex-Levenberg-Marquardt algorithm was used in the minimization of the objective function. The numerical algorithms were included in the program package MODEST.24 The minimization was commenced by the simplex algorithm but was switched to the more rapid Levenberg-Marquardt algorithms as the minimum was approached. The model’s degree of explanation is defined by eq 7, in which cjA is the average value of the experimentally recorded concentrations.

R2 ) 1 -

∑(cA(t) - cˆ A(t))2 ∑(cA(t) - cjA)2

Figure 8. [OCN-] vs time; O experimental data from experiment 2, - model calculations from parameter estimation.

(7)

Parameter estimation of our experiments 1-4 gave similar values for k3 and kHCO3- as Jensen’s runs 5-714 and run 5012 (Table 4). This is reasonable, as our experiments 1-4 were replicates of the given runs by Jensen. However, when all of our experimental data from the higher concentrations of hydroxide and carbonate were included, the value for k3 was similar, but the value for kHCO3- was significantly higher. Fixing p1 (k3) in the parameter estimation lowered the value for kHCO3- , based on all experimental data from this work, to a degree, but it is still higher than for the solutions of lower ionic strength. This discrepancy in the value of kHCO3- is likely due to the calculated ratio of activity coefficients, fCO32-/fOH- fHCO3- , being too low for the solutions of higher ionic strength. Figures 8 and 9 are data fits from parameter estimation, in which p1 is fixed to 1.70 × 10-3. The degree of explanation is 99.4 and 94.9%, respectively. Temperature Dependence of Rate Constant. Table 5 gives the rate constants k3 and kHCO3- at various temperatures from both literature and this work. From the literature findings and our experiments 1-4 it is possible to calculate the temperature dependence of these rate constants using the Arrhenius equation, leading to eqs 8 and 9.

Figure 9. [OCN-] vs time; O experimental data from experiment 2, - model calculations from parameter estimation.

-11600 (min-1), R2 ) 0.997 (8) T

k3 ) 7.96 × 1010e

-8725 (L min-1 mol-1), T R2 ) 0.997 (9)

kHCO3- ) 4.55 × 109e

The activation energies for k3 and kHCO3- were calculated to be 96 and 73 kJ/mol, respectively. Predicted vs Experimental Cyanate Decomposition in Three Mill Green Liquors. Kinetic experiments were carried out at 95 °C with green liquors from the smelt dissolving tank of three kraft pulp mills. The activity coefficients were calculated as for the other

4820 Ind. Eng. Chem. Res., Vol. 43, No. 16, 2004

Figure 10. Concentration of OCN- vs time for the three mill liquors at 95 °C. Experimental results: [ mill A; ∆ mill B; O mill C; model predictions: - mill A; - - - mill B; - - - mill C. Table 5. Values for Rate Constants from Literature and This Study refs 12, 14

13 this study expt 1-4 this study all expt

k3 (min-1)

kHCO3- (L mol-1 min-1)

1.7 × 10-4 (70 °C) 5.7 × 10-4 (82 °C) 1.7 × 10-3 (94 °C)

4.8 × 10-4 (18 °C) 9.2 × 10-4 (25 °C) 6.6 × 10-3 (50 °C) 0.045 (70 °C) 0.21 (94 °C)

10-6 (30

1.8 × °C) 1.9 × 10-3 (100 °C) -3 1.7 × 10 (94 °C) 1.6 × 10-3 (94 °C)

0.24 (94 °C) 0.37 (94 °C)

solutions and were given in Table 3. Eqs 2, 8, and 9 were used to predict the decomposition of cyanate with time for the three different green liquors and is compared to the experimental results in Figure 10. The degree of explanation is 87.2, 94.6, and 81.6% for mill liquors A, B, and C, respectively. The error bars for the mill C experimental data represents one standard deviation based on six replicate analysis for each sample. The experimental data for mills A and B were analyzed two to three times each. Due to the high dilution required for analysis of the mill liquors and the low initial cyanate concentration, analysis became a limiting factor in the experiments with green liquor, and, therefore, the runs were of shorter duration that the runs with the model solutions. As clearly follows from Figure 10, the structure of the kinetic model developed based on experiments with model mixtures as well as the calculated values of kinetic constants was capable of predicting the decomposition of cyanate in alkaline solutions of the tested liquors obtained from several kraft pulp mills. Conclusions The experiments with varying concentrations of carbonate and hydroxide show that bicarbonate is the catalytically active species in cyanate decomposition, even in highly alkaline conditions where its concentration is very low. The concentration of bicarbonate in these solutions can be determined using acid-base equilibrium, but activity must be accounted for at these pH levels, where even at fairly low concentrations of carbonate, there is a high ionic concentration. It appears possible to use the more easily calculated ionic strength to estimate the ratio of the activity coefficients in eq 3 when it is not possible to make the equilibrium calcula-

tions. Experiments varying the ionic strength with other anions that do not take part in the reaction should be performed to confirm this. The parameter estimation method was used to determine the values of the rate constants k3 and kHCO3- for the derived rate equations. The rate constant k3 was equivalent to that found by Jensen for identical experiments, while it was 6% lower when all of our experimental data were considered. The rate constant, kHCO3-, was 14% higher for identical experiments and 57% higher when all experimental data were considered. The larger discrepancy when taking into account all of the data is most likely due to the calculated ratio of activity coefficients being to low. Our results were coupled with those from literature to obtain an Arrhenius description of the temperature dependency of the rate constants k3 and kHCO3-. The resulting kinetic expression allows the rate of cyanate decomposition (ammonia formation) to be calculated in highly alkaline solutions of high ionic strength such as those of the kraft pulping chemical recovery process. This method of combining kinetic and thermodynamic calculations allows for an improved description of industrial solutions. Acknowledgment We thank Zhang Shengqiang for his experimental work, Linus Perander for performing both experimental work and parameter estimation calculations, and Luis Bezerra for his work in cyanate analysis. The work with the prepared aqueous solutions was done under the PCG Anions project funded by the Academy of Finland as a part of its Center of Excellence Program. The work done with the kraft pulp mill green liquors was completed under the EU REMPULP project “Reduction of Air Emissions at Kraft Pulp Mills” (contract QLK51999-01105). We would like to acknowledge the support of Andritz Oy, Foster Wheeler Energia Oy, Kvaerner Pulping Oy Power Division, Oy Metsa¨-Botnia Ab, and Vattenfall Utveckling AB and the Finnish National Technology Agency (TEKES) in the Chemistry in Biomass Combustion project (ChemCom). The stipend for the graduate studies of DeMartini from the Graduate School in Chemical Engineering is gratefully acknowledged. Literature Cited (1) Kyma¨la¨inen, M.; Forsse´n, M.; Hupa, M. The Fate of Nitrogen in the Chemical Recovery Process in a Kraft Pulp Mill Part I. A General View. J. Pulp Paper Sci. 1999, 24(12), 410. (2) DeMartini, N. Determination of Cyanate in Kraft Green Liquor. M.Sc. Thesis. Åbo Akademi University, Turku, Finland, 2001. (3) Kyma¨la¨inen, M.; Forsse´n, M.; DeMartini, N.; Hupa, M. The Fate of Nitrogen in the Chemical Recovery Process in a Kraft Pulp Mill Part II. Ammonia Formation in Green Liquor. J. Pulp Paper Sci. 2001, 27(3), 75. (4) DeMartini, N.; Forsse´n, M.; Niemela¨, K.; Samuelsson, A° .; Hupa M. Release of Harmful Nitrogen Species from the Recovery Processes of Three Kraft Pulp Mills. In Proceeding to the 2003 TAPPI Fall Technical Conference: Engineering, Pulping and PCE&I, Oct. 26-30, Chicago, IL. (5) NCASI. Effect of Stripper Off-Gas Burning on NOx Emissions. Technical Bulletin no. 802. NCASI: New York, March 2000. (6) Veverka, P.; Nichols, K.; Horton, R.; Adams, T. In 1993 TAPPI Proceedings Environmental Conference; TAPPI: Boston, MA, 1993, p 777. (7) Niemela¨ K.; Ulmgren, P. Behaviour of Nitrogen During Kraft Pulping of Wood. In The 7th European Workshop on Lignocellulosics and Pulp Oral Presentaitons, Turku, Finland, 2002, p 71.

Ind. Eng. Chem. Res., Vol. 43, No. 16, 2004 4821 (8) Kyma¨la¨inen, M.; Forsse´n, M.; Jansson, M.; Hupa, M. The Fate of Nitrogen in the Chemical Recovery Process in a Kraft Pulp Mill. Part IV. Smelt Nitrogen and Its Formation in Black Liquor Combustion. J. Pulp Paper Sci. 2002, 28(5), 151. (9) Sidgwick, N. V. The Chemical Elements and Their Compounds; Clarendon: Oxford, 1950; p 673. (10) Lister, M. W. Some Observations on Cyanic Acid and Cyanates. Can. J. Chem. 1955, 33, 426. (11) Kemp, I. A.; Kohnstam, G. The Decomposition of Inorganic Cyanates in Water. J. Chem. Soc. 1956, 900. (12) Jensen, M. B. On the Kinetics of the Decomposition of Cyanic Acid. Acta Chem. Scand. 1958, 12, 1657. (13) Vogels, G. D.; Uffink, L.; Van Der Drift, C. Cyanate Decomposition Catalyzed by Certain Bivalent Anions. RECUEIL 1970, 89, 500. (14) Jensen, M. B. On the Kinetics of the Decomposition of Cyanic Acid II: The Carbonate Catalysis. Acta Chem. Scand. 1959, 13, 659. (15) O’Neil, M. J.; Smith, A.; Heckelman, P. E.; et al. The Merck Index; Merck and Co.: New Jersey, 2001; p 306. (16) DeMartini, N.; Kyma¨la¨inen, M.; Forsse´n, M.; Hupa, M. Quantification of Cyanate in Pulp Mill Recovery Streams Using Ion Chromatography. In 2001 TAPPI International Chemical Recovery Conference Poster Presentations, June 11-14, Whistler, BC, Canada, p 83.

(17) Ha¨gg, G. Kemisk Reaktionsla¨ ra; Almqvist & WiksellGeber: Uppsala, 1954. (18) Sille´n, L. G. In Treatise on Analytical Chemistry Part 1: Theory and Practice; Kolthoff, Elvin, G., Eds.; Vol. I, Interscience Encyclopedia: New York, 1959; Chapter 8. (19) Ericksson, G.; Hack, K. ChemSage - A Computer Program for the Calculation of Complex Chemical Equilibria. Metallur. Trans. B 1990, 21B, 1013. (20) Pitzer, K. S. Activity Coefficients in Electrolyte Solutions; CRC Press: Boston, MA, 1991; 542p. (21) Bale, V. W.; Pelton, A. D.; Thompson, W. T. The FACT Database; Ecole Polytechnique de Montreal/Royal Military College: Canada, 1997. (22) Taylor, C. E. Thermodynamics of Sodium Carbonate in Solution. J. Phys. Chem. 1955, 59, 653. (23) Harned, H. S. The Activity Coefficient of Sodium Hydroxide in Aqueous Solution. J. Am. Chem. Soc. 1925, 47, 676. (24) Haario, H. Modest 6.0 - A User Guide; ProfMath: Helsinki, 2001.

Received for review October 16, 2003 Revised manuscript received April 30, 2004 Accepted May 23, 2004 IE034189+