Effects of Temperature and Pressure on the Thermolysis of Morpholine

Feb 18, 2015 - Amines have been applied in nuclear power plants for decades, and some guidelines are available.(3) However, the use of ... Retention t...
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Effects of Temperature and Pressure on the Thermolysis of Morpholine, Ethanolamine, Cyclohexylamine, Dimethylamine, and 3‑Methoxypropylamine in Superheated Steam David H. Moed,*,† Arne R. D. Verliefde,‡ and Luuk C. Rietveld† †

Faculty of Civil Engineering and Geosciences, Delft University of Technology, Stevinweg 1, 2628 CN Delft, The Netherlands Particle and Interfacial Technology Group, Ghent University, Coupure Links 653, 9000 Gent, Belgium

‡

S Supporting Information *

ABSTRACT: Alkalizing amines such as cyclohexylamine and dimethylamine have great potential for protecting steam−water cycles against corrosion, but their thermal stability is limited and anionic decomposition products are a concern because of increased corrosion risk. In this study, morpholine, ethanolamine, cyclohexylamine, dimethylamine, and 3-methoxypropylamine were exposed to temperatures of 500, 530, and 560 °C at pressures of 9.5, 13.5, and 17.5 MPa to investigate the effects of temperature and pressure on amine thermolysis kinetics. The surface/volume ratio of the reactor tube was 0.4 mm−1, close to the value for superheater tubes in steam−water cycles. All amines thermolyzed by first-order kinetics, with the exception of dimethylamine. The Arrhenius constants Ea, ln(A), and Va were obtained from the experimental data for all investigated amines. The effect of pressure on the thermolysis kinetics was less pronounced than in previous studies and was different for each amine. Dimethylamine did not degrade below 20% and 10% at 500 and 530 °C, respectively, despite the application of longer retention times, suggesting the synthesis might occur. Limited practical data showed some promise for the applicability of the model to steam−water cycles. More plant data are needed to fully validate the model. In all cases, thermolysis of the amines led to the formation of between 150 and 600 ppb organic acid anions. In most cases, the concentrations increased linearly with increasing degradation percentage. Acetate and formate were found to be major degradation products, with some propionate and traces of glycolate. Cationic degradation products were ammonia and some amines, meaning that the complete thermolysis of an amine does not necessarily lead to acidic conditions.

1. INTRODUCTION The demineralized water used in steam−water cycles is chemically conditioned to raise the pH of the water and reduce failures due to corrosion. One of the applied chemical treatment methods is all-volatile treatment (AVT), which uses ammonia to raise the pH of both the boiler (feed) water and the steam or condensate. Although AVT is commonly applied for steam−water cycles (SWCs), AVT requires high water purity1 and still enhances the corrosion of any copper that might be present in an SWC. Also, because ammonia is more volatile than water, it offers limited protection to corrosion in areas where flashing occurs or the first droplets of condensate are formed.2,3 Alkalizing amines such as morpholine (MOR) and ethanolamine (ETA) offer a less volatile alternative to ammonia,4 making the SWC more resilient to corrosion.5−7 Because ammonia corrodes copper, amines can also be a good chemical water treatment solution for systems with mixed metallurgy. Amines have been applied in nuclear power plants for decades, and some guidelines are available.3 However, the use of amines in fossil-fired plants is under debate, because of the higher operating temperatures and pressures employed. Because amines are known to have limited thermal stability, they can decompose to form organic acid anions, lowering pH and thereby increasing corrosion risks.2,8 Potential formation of organic acid anions during amine thermolysis is currently preventing the general acceptance of amines for application in steam−water cycles. © 2015 American Chemical Society

Some experimental data on the thermal degradation of amines are available, but most such data focus only on conditions found in nuclear plants,9−12 with data applicable to fossil-fired plants being more limited. Mori et al.13 tried to simulate the thermolysis of six amines under fossil-fired plant conditions at higher temperatures with a boiler−superheater test loop. They considered the effect of the superheater on thermolysis by varying the superheater exhaust temperature. They concluded that MOR breakdown in the boiler was undetectable and started only from 500 °C onward in the steam phase, whereas ETA started degrading already at steam exhaust temperatures of 450 °C. The retention time was not mentioned, making any calculation of kinetics impossible. The importance of the superheater in the applicability of alkalizing amines in fossil-fired plants was also pointed out by Bull14 in a report detailing the results from a fossil-fired plant survey on the use of amines. In another study, MOR was dosed (9 ppm) to investigate the effects of pressure, temperature, and retention time on amine degradation kinetics in a flow reactor, mimicking a superheater.15 MOR degradation kinetics were found to be firstorder, increasing with temperature and decreasing with Received: Revised: Accepted: Published: 2606

December 12, 2014 February 13, 2015 February 18, 2015 February 18, 2015 DOI: 10.1021/ie504849v Ind. Eng. Chem. Res. 2015, 54, 2606−2612

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Figure 1. Schematic of the flow reactor used during the experiments (from left to right): argon, bottle containing the solution of interest, HPLC pump, fluidized sand bed (FSB), metal tube in the FSB, cooling spiral, and back-pressure regulator (BPR).

pressure. The tubes used in the flow reactor, however, were only 1.5 mm in diameter, making the surface/volume (S/V) ratio (4.65 mm−1) much higher than in full-scale fossil-fired plants, where the tube diameters are typically at least 10 times larger. The metal oxide layer on the surface of the tubes appeared to also influence the thermolysis kinetics. In a previous publication, we provided a relation between the S/V ratio and thermolysis kinetics and made a distinction between homogeneous (in the bulk steam) and heterogeneous (on the metal surface) thermolysis.16 Although heterogeneous thermolysis accounted for 82−92% of the value of the degradation rate constant k for tubing with an S/V ratio of 4.65 mm−1, the contribution was only 6−17% at an S/V ratio of 0.4 mm−1. This led to the conclusion that an amine thermolysis study is best conducted with larger metal tubes, preferably with S/V ratios of 0.4 mm−1 or less, to approximate superheater tube sizes. In this study, the effects of temperature and pressure on the thermolysis of five amines that are being applied in steam− water cycles [MOR, ETA, dimethylamine (DMA), cyclohexylamine (CHA), and 3-methoxypropylamine (MOPA)] were studied in a flow reactor equipped with a Hastelloy C-276 tube with an S/V ratio of 0.4 mm−1. The data were used to construct a predictive model to assess the applicability of amines in hightemperature and -pressure steam−water cycles. The organic acid anions and ammonia that are produced during amine thermolysis are discussed, and an explanation is given as to why DMA does not degrade according to first-order kinetics.

Table 1. The Investigated Amines, Temperatures, Pressures, and Retention Timesa amineb

T (°C)

P (MPa)

t (s)

MOR ETA DMA CHA MOPA

500 530 560

9.5 13.5 17.5

0−5 5−10 10−20 20−30

a

T, temperature; P, pressure; t, retention time of the steam inside the tube. bEach amine has been tested at every temperature, pressure, and retention time listed.

argon, and 0.1 ppm carbohydrazide was dosed to eliminate any remaining traces of oxygen. A more detailed description of the setup and procedure was already published by Moed et al.16 2.2. Analyses. Analysis of organic acid anions was performed using a Metrohm (Schiedam, The Netherlands) 881 ion chromatography system. A Metrohm A Supp 16 4.0/ 250 anion column was operated at 67 °C with eluent containing 7.5 mM Na2CO3 + 0.75 mM NaOH in ultrapure water. The suppressor was regenerated with 50 mM H2SO4, and the limit of detection was 1 ppb. For cation analysis, a Metrohm C5 cation column was used, with 3 mM HNO3 as the eluent, for which the limit of detection for MOR, ETA, MOPA, and DMA was 5 ppb and that for CHA was 50 ppb. The detection limit for ammonia was 1 ppb. 2.3. Temperature and Pressure Relations. First-order isobaric degradation kinetics as a function of absolute temperature were modeled according to the equations17

2. MATERIALS AND METHODS 2.1. Setup and Conditions. A schematic of the experimental setup is provided in Figure 1. A high-performance liquid chromatography (HPLC) pump sends the solution of interest through Hastelloy C-276 tubing inside a fluidized sand bath at a fixed flow rate. The tube inside the fluidized sand bed was a 12.5-mm Hastelloy C-276 tube with a volume of 6.2 mL and an internal diameter of 10 mm, thus having an S/V ratio of 0.4 mm−1. The amine solutions of interest were tested under the conditions listed in Table 1, and the experiments were executed in duplicate. Retention times were varied by adjusting the HPLC pump flow rate to 0.75, 1.0, 1.5, or 2.5 mL/min, and the flow rate was checked for accuracy before every run. Because water and steam increase in volume at higher temperatures, the retention time in the tubing is a function of flow rate, temperature, and pressure inside the tube. Therefore, applied retention times are listed in ranges. Oxygen was removed with

r = k(T )C

(1)

k(T ) = A e−Ea / RT

(2)

where r is the degradation rate, k(T) is the degradation rate constant as a function of temperature, C is the concentration of the degrading compound, A is the pre-exponential factor, Ea is the activation energy, T is the absolute temperature, and R is the universal gas constant. The latter equation can be written in logarithmic form as ln(k) = ln(A) −

Ea 1 RT

(3)

Therefore, in an Arrhenius plot of ln(k) versus 1/T, the slope of the regression line through the data points gives the experimental value of Ea/R, and the y intercept corresponds to ln(A). 2607

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Industrial & Engineering Chemistry Research In a previous study, Moed et al.15 showed the independency of the activation energy (Ea) and activation volume (Va) on both temperature and pressure. When the natural logarithm of the degradation rate constant (k) is plotted as a function of pressure (at a constant temperature), the activation volume can be derived from the slope ln(k) = ln(k ref ) −

Va (P − Pref ) RTref

steam. The linear regression lines given for each pressure enabled the calculation of the activation energy Ea (by using the slope) and the prefactor A (by using the y intercept). The averages of the resulting three Ea and ln(A) values were used in the model. The same data were used to plot ln(k) against P for each temperature (see Figure 3), from which Va was calculated

(4)

where kref is the degradation rate constant corresponding to the reference temperature (Tref) and reference pressure (Pref). The calculated values for Va should be independent of the chosen values for Tref and Pref. Combining eq 3 and 4, the reaction rate as a function of temperature and pressure can be written as ⎤ ⎡ E V r = k(T , P)C = exp⎢ln(A) − a − a (P − Pref )⎥C ⎦ ⎣ RT RT (5)

To determine the experimental values for k, an integrated first-order rate law was fitted to the experimentally determined decrease in concentration of the amines upon degradation. First order was assumed because it was found to be the case for amine degradation in most previous studies.9,10,12,15,16,18,19 C = C0e−kt

Figure 3. Arrhenius plot of the relation between ln(k) and P for the thermolysis of ETA in dry steam at 9.5, 13.5, and 17.5 MPa.

using the slope of the linear regression lines at each temperature. The average for the three temperatures was the value of Va used in the proposed model. The Arrhenius plots for the other four amines can be found in the Supporting Information. Figures S5, S8, and S11 show the pressure relation of ln(k) for MOR, MOPA, and CHA and have almost horizontal regression lines. Figures S4, S7, and S10 (Supporting Information) show their temperature relation of ln(k), where the regression lines are close together. MOR, MOPA, and CHA therefore show no clear pressure dependency. All values found for Ea, A, and Va are summarized in Table 2, including the errors. The activation volumes of MOR, MOPA, and CHA have been highlighted in italics, because of their apparent lack of pressure dependency. It must be noted that a higher activation energy does not necessarily imply higher thermal stability of a compound, as thermal stability also depends on the prefactor A. The activation energy indicates the amount of energy that is

(6)

The reaction rate constant k can be found by plotting C/C0 against time and calculating the exponential regression function y = e−kx. The threshold for a good fit was chosen at R2 = 0.95, and when the results that did not meet this requirement, the corresponding experiments were repeated. Subsequently, A, Ea, and Va can be determined from eqs 3 and 4 as described above.

3. RESULTS AND DISCUSSION 3.1. Temperature and Pressure Relations. Degradation curves were constructed from the obtained data. The thermolysis of all amines led to first-order regression curves, with the exception of DMA, which will be discussed in more detail later. First-order kinetics applied to degradation percentages over 90% as well, showing concentration independency in the range of 1−9 ppm. This range could be larger, but the data cannot validate such a claim. The values for k resulting from the degradation curves were used to create Arrhenius plots for the temperature relation. An example of such a plot is given in Figure 2, which shows the relation between ln(k) and 1/T for the thermolysis of ETA in dry

Table 2. Summary of Ea, ln(A), and Va Values Found for Each Amine, Including Standard Deviations (SDs) and Relative Standard Deviations (RSDs) amine ETA

MOR

MOPA

CHA

DMA

Figure 2. Arrhenius plot of the relation between ln(k) and 1/T for the thermolysis of ETA in dry steam at 500, 530, and 560 °C. 2608

constant

mean

SD

RSD (%)

units

Ea ln(A) Va Ea ln(A) Va Ea ln(A) Va Ea ln(A) Va Ea ln(A) Va

189 25.5 369 295 40.2 137 254 35.6 6 409 56.9 272 100 13.6 346

1 0.4 59 15 2.2 134 16 2.4 151 12 1.7 112 3 0.4 19

1 2 16 5 5 98 6 7 2343 3 3 41 3 3 6

kJ mol−1 s−1 cm3 mol−1 kJ mol−1 s−1 cm3 mol−1 kJ mol−1 s−1 cm3 mol−1 kJ mol−1 s−1 cm3 mol−1 kJ mol−1 s−1 cm3 mol−1

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500 °C, thermal degradation stopped when DMA was at around 20%, as shown in Figure 5B. At 530 °C, the DMA concentration did not drop below 10%. At 560 °C, the DMA concentration went down to less than 2% within a few seconds, but the compound never fully degraded. In some cases, DMA concentrations measured for longer retention times were even higher than those measured for shorter retention times, and a clear increasing trend was observed. For instance, at 500 °C and 17.5 MPa, 15% DMA remained after 14 s in the reactor. After 43 s under the same conditions, 23% DMA was left. This increase cannot be explained through the error of detection of the ion chromatography system. Longer retention times do not allow degradation of DMA below 20% at 500 °C, suggesting that the DMA thermolysis reaction is reversible. This can happen only if DMA is synthesized from its breakdown products. Mochida et al.20 investigated the synthesis of DMA and methylamine (MET) from ammonia and methanol over metal oxides at 300−450 °C, which could yield up to 50%. A patent by Ashina et al.21 describes almost the same process, but claims that higher selectivity for DMA can be obtained by exposing methanol and ammonia to steam at 250−700 °C. During the experiments, ammonia was a degradation product, and the presence of methanol was plausible. MET was also found as a degradation product in all DMA experiments, and its concentration remained constant at higher retention times. Although this reasoning is speculative, the synthesis of DMA in a superheater, or perhaps in the condensing stages, is plausible. As mentioned before, DMA was much more stable in a 1.6mm tube at 490 °C18 than in a 12.5-mm tube at 500 °C. This indicates that the synthesis of DMA benefits from a high S/V ratio. Having relatively more metal oxide surface area would benefit the synthesis of DMA, thereby increasing the apparent DMA stability as the S/V ratio increases. 3.3. Organic Acid Anion Formation. Potential formation of organic acid anions during amine thermolysis is currently preventing the general acceptance of amines for application in steam−water cycles. In a previous amine thermolysis study,16 formate and acetate were found to be the major anionic degradation products of MOR and ETA. In addition, glycolate was measured as a degradation product of MOR and ETA. In the current study, however, propionate was also found in samples from MOR degradation. All organic acid data from MOR thermolysis experiments conducted in the current study are summarized in Figure 6. Formate was the major degradation product, which does not correspond to the findings in the previous study with a higher S/V ratio. This difference indicates that the surface of the tubes has an affect not only on the degradation kinetics, but also on what degradation product is preferably formed. Formate, acetate, and propionate all increased with increasing degradation (despite the increasing temperature), whereas glycolate was found at only trace levels throughout all MOR thermolysis experiments. The other organic acid plots for all of the investigated amines can be found in the Supporting Information. It can be seen that formate and acetate from ETA degradation were found at similar concentrations (by mass) and their concentrations increased linearly to 290 ppb for 100% amine degradation. The glycolate concentration after ETA degradation seemed to decrease with increasing ETA degradation, but also decreased with increasing temperature, which indicates the low thermal stability of glycolate. For MOPA, the highest concentrations found for the degradation products were those of formate (up

required to start the degradation reaction, but the prefactor A determines how quickly the amine degrades as a function of temperature. For example, the thermolysis of CHA in dry steam has a high activation energy at 500 °C, resulting in a higher value for k than for MOR, making CHA more stable than MOR at that temperature. However, the high prefactor determines that the k values of CHA and MOR are almost the same at 560 °C and 13.5 MPa (0.074 and 0.072 s−1, respectively). A previous study showed much lower MOR stability than observed here.15 That study was executed with a smaller stainless steel tube (S/V = 4.65 mm−1), resulting in a large contribution of catalytic wall effects to the thermolysis. This emphasizes the importance of the S/V ratio when studying thermal degradation under conditions resembling those in fullscale plants, as was shown in ref 16. In the previous study, a much greater effect of pressure for MOR was found as well. This suggests that metal surface catalysis enhances the pressure dependency effect. Another report discussed the thermal stability of MOR, CHA, ETA, and DMA in an experimental setup similar to that used in this work but with an S/V ratio of 4.65 mm−1.18 MOR, CHA, and ETA showed lower thermal stability in the smaller tubing. ETA, for instance, underwent 90% degradation in 12 s at only 490° and 17.5 MPa. DMA was an exception, showing higher stability in the small tubing. DMA thermolysis, however, did not always follow first-order kinetics, thus making modeling difficult, as discussed in the next section. To compare the experimentally obtained values of k with those calculated using eq 5 and the results in Table 2, the experimental k values were plotted against the calculated k values. The coefficient of determination R2 was used as a measure for the goodness of fit. As an example, the relation between the measured and calculated values of k for ETA is shown in Figure 4. The figures for all other amines are provided

Figure 4. Calculated values of k plotted against the experimentally obtained values of k for all experiments conducted with ETA.

in the Supporting Information. The coefficients of determination of the modeled compared to the experimental data for ETA, MOR, MOPA, CHA, and DMA were found to be 0.99, 0.97, 0.97, 0.92, and 0.99, respectively. 3.2. Dimethylamine. DMA requires special attention among the studied amines because it was the only compound that did not degrade by first-order kinetics in all situations. The previously discussed Arrhenius plots and model calculations for DMA were made with k values derived from the shortest retention times (Figure 5A), but when DMA was exposed to longer retention times, first-order kinetics no longer applied. At 2609

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Figure 5. Degradation percentage vs retention time for the thermolysis of DMA at 500 °C for (A) shorter and (B) longer retention times. First-order degradation does not apply to longer retention times.

meaning that a large proportion of the amines degraded to neutral compounds. In all cases, thermolysis of the amines led to the formation of organic acid anions. In most cases, the concentrations increased linearly with increasing degradation percentage. Organic acid anions are a concern in power-plant regions with two-phase flow, because organic acids are more soluble than ammonia. This means that, when amines are used, there might be a point at which the acidity of the degradation products has enough effect on the water phase that the solubility of magnetite (the protective oxide layer found under power-plant conditions) increases. In practice, not all amines would be dosed at the same concentration, as their dosages will depend on their basicities. An amine that requires a lower dose to maintain sufficiently high pH in a steam−water cycle will also form fewer organic acid anions. It also has to be noted that amines and organic acid anions in a power plant will cycle through the system, so the thermal stability of organic acids in superheated steam matters as well. Regardless, the protection an amine offers is a combination of its thermal stability and its tendency to produce organic acid anions. 3.4. Nitrogen Balance. As an amine breaks down, the cationic degradation products could still offer some protection. Ammonia and other cationic degradation products were measured in this study to determine the percentage of nitrogen that can be found as cations providing alkalinity. This quantity is expressed as a percentage of the total organic nitrogen in the influent. For MOR, which degraded into ammonia and small traces of ETA (up to 229 ppb), the results for the cationic degradation products are shown in Figure 7. Not all nitrogen

Figure 6. Acetate, formate, glycolate, and propionate formation vs degradation percentage during thermolysis of MOR at 500, 530, and 560 °C and 9.5, 13.5, and 17.5 MPa.

to 242 ppb) and acetate (up to 120 ppb). The propionate concentration increased slightly with increasing MOPA degradation, but did not exceed 17 ppb. CHA produced only a low amount of organic acids, with the highest total organic acid anion concentration being less than 200 ppb. The acetate concentrations from CHA were slightly higher than the formate concentrations, which were slightly higher than the propionate concentrations. All three organic acid anions increased proportional to CHA degradation. An offset of 5 ppb formate and 10 ppb acetate can be seen for low degradation percentages. This is caused by organic carbon contamination (from, e.g., the air or the interior of the plastic pump), which is hard to prevent, even in full-scale plants. This is also the reason that DMA thermolysis seemed to produce small amounts of acetate, which was not the case. DMA produced almost exclusively formate, which can be attributed to its small molecule structure, making the formation of other organic acid anions improbable, unless synthesis occurred. Formate was also observed to be the only anionic degradation product by Bull14 when using DMA as an ammonia replacement for a secondary circuit. Although the acetate and propionate concentrations did not fluctuate, the formate concentrations went up and down for all amines, even under similar conditions and degradation percentages. This can be seen mostly for ETA, MOPA, and DMA, where the formic acid concentrations for high degradation percentages varued in the ranges 159−271, 51− 242, and 62−299 ppb, respectively. What causes the formate concentrations to be so unpredictable remains unclear. Over the entire study, 1−15% (by mass) of the carbon released by degraded amines could be traced back as organic acid anions,

Figure 7. Percentage of total nitrogen measured as cationic degradation products (theoretical and measured) for MOR at 500, 530, and 560 °C and 9.5, 13.5, and 17.5 MPa. 2610

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A difference in model and plant data was found in the formation of organic acids in plant A, where ETA was dosed. Almost no formate was measured at this location, which is different from what was measured in plant B and the observations during the ETA degradation in this study. Plant B had a formate/acetate ratio close to what was experimentally obtained in this laboratory-scale study. In plant C, the organic acid concentrations were very low, but MOPA was dosed at only 20 ppb in this case.

could be traced back as ammonia or other cationic degradation products. This is probably due to the degradation of the amines and/or ammonia to inorganic nitrogen or to the loss of ammonia (because of its volatility) during grab sampling. A blank run with only 0.1 ppm carbohydrazide revealed that 33 (±15) ppb ammonia came from the carbohydrazide. The results were corrected for this contribution. For other amines, a higher degree of scattering can be seen in the nitrogen recovery plots, which are shown in the Supporting Information. In some cases, 100% of the nitrogen can be accounted for, usually at lower temperatures or lower degradation percentages. In all cases, the dominant cationic degradation product was ammonia. For DMA, traces of MET (up to 236 ppb) were found as a decomposition products as well, whereas MOPA also degraded to some hydroxyl propylamine (up to 204 ppb). With most nitrogen always being converted into ammonia or another amine, the thermolysis of an amine in a power plant does not necessarily have to lead to a hazardous pH drop, as the pH effect of the formation of organic acids will mainly be countered by the alkalizing effect of the ammonia. As stated in the previous section, this could still imply a lower degree of protection in two-phase flow, as some organic acids are less volatile than ammonia and prefer the liquid film that covers the metal surface. 3.5. Comparison with Practical Data. Practical data in the literature that can be used to verify the model are almost nonexistent, mostly because of the lack of information on retention times in superheaters and reheaters. Some general comparisons can be made, however. Layton and Daniels22 reported a power-plant survey on amine usage that included detailed data for ETA degradation in several parts of two plants (A and B), and in plant B, CHA was also dosed. The third plant (C) ran (partially) on MOPA dosing. The superheat (SH) and reheat (RH) temperatures of plants A−C were 540, 566, and 538 °C, respectively, with pressures of 17.9, 13.1, and 13.1 MPa, respectively. Retention times could not be calculated from the reported data, so degradation percentages can be compared only relative to each other. Similarities were found between their study and this study in the stability of CHA relative to ETA. Layton and Daniels22 showed a loss of CHA and ETA over the SH and RH, but also showed that CHA is more stable than ETA at 551 °C. The average loss percentages of CHA and ETA were (47 ± 7)% and (60 ± 11)%, respectively, at a 551 °C RH steam temperature. According to the model presented in this article, a 47% decrease in CHA at 551 °C and a pressure of 13.1 MPa corresponds to a 70% loss of ETA under similar conditions. As such, the findings corroborate the practical data (within the degree of accuracy of the plant data). When applying the same retention time to the conditions found in plant C (a heat recovery steam generator similar to plant B), the model data predict a 70% loss of MOPA, which is close to the highest degradation percentage of MOPA found in that specific plant of 66%. It should be kept in mind that comparisons between model data based on laboratory-scale experiments and full-scale data need to be handled with care. Amines in a plant do not stay at a fixed temperature for a few seconds, but follow a profile of temperatures. Determining retention times and temperature profiles in superheaters and reheaters is a challenge. However, making these kinds of comparisons is important for assessing the applicability of amine thermal stability modeling.

4. CONCLUSIONS This study has led to the following conclusions: (1) All amines thermolyzed by first-order kinetics, with the exception of dimethylamine. (2) The obtained Arrhenius constants are summarized in Table 3. Table 3. Summary of Obtained Arrhenius Constants ETA MOR MOPA CHA DMA

Ea (kJ mol−1)

ln(A) (s−1)

Va (cm3 mol−1)

189 295 254 409 100

25.5 40.2 35.6 56.9 13.6

369 137 6 272 346

(3) The effect of pressure on the thermolysis kinetics was much less pronounced than in previous studies, which can probably be attributed to the low S/V ratio in this study, which was applied to come closer to superheater conditions. (4) Dimethylamine did not degrade below 20% and 10% of the initial concentration at 500 and 530 °C, respectively, despite the application of longer retention times. This suggests that synthesis might occur, either at high temperature or in the condensing stages. Further studies are needed to clarify this phenomenon. (5) In all cases, thermolysis of the amines led to the formation of organic acid anions. In most cases, the concentrations increased linearly with increasing degradation percentage. Acetate and formate were found to be the major degradation products, with some propionate and traces of glycolate. (6) Cationic degradation products consisted of mostly ammonia and some amines, meaning that the complete thermolysis of an amine does not necessarily lead to acidic conditions. Not all nitrogen could be traced back to these cationic degradation products, suggesting that some other inorganic nitrogen species were formed or that ammonia was lost during sampling. (7) Comparison of the model to some limited practical measurements of amines in power plants shows some promise in the applicability of the model. A larger data set and more information on superheater retention times are necessary to validate the model proposed in this article, which should try to focus on retention times in the hottest sections of power plants.

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ASSOCIATED CONTENT

S Supporting Information *

Arrhenius, organic acid anion, and nitrogen recovery plots are provided to support the findings described in this article. This material is available free of charge via the Internet at http:// pubs.acs.org. 2611

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International Conference on the Properties of Water and Steam, London, U.K., Sep 1−5, 2013. (19) Moed, D. H.; Verliefde, A. R. D.; Rietveld, L. C.; Heijman, S. G. J. Organic acid formation in steam-water cycles: Influence of temperature, retention time, heating rate and O2. Appl. Therm. Eng. 2014, 65, 194−200. (20) Mochida, I.; Yasutake, A.; Fujitsu, H.; Takeshita, K. Selective synthesis of dimethylamine (DMA) from methanol and ammonia over zeolites. J. Catal. 1983, 82, 313−321. (21) Ashina, Y.; Fujita, T.; Fukatsu, M.; Yagi, J. Process for producing dimethylamine in preference to mono- and trimethylamines by gas phase catalytic reaction of ammonia with methanol. U.S. Patent 4,582,936, 1986. (22) Layton, K. F.; Daniels, D. G. Interim GuidanceAmine Treatments in Fossil Power Plants; EPRI: Palo Alto, CA, 2010.

AUTHOR INFORMATION

Corresponding Author

*E-mail: [email protected]. Tel.: +31-0-152786588. Notes

The authors declare no competing financial interest.

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ACKNOWLEDGMENTS The authors thank EPRI and Evides Industriewater for financial and technical support. Specific gratitude goes to Michael Caravaggio and James Mathews from EPRI for useful input and discussions.

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REFERENCES

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DOI: 10.1021/ie504849v Ind. Eng. Chem. Res. 2015, 54, 2606−2612