# New insight on the impact of NOx on diethanolamine (DEA) degradation mechanisms.

Mohammad Shoeb, Girel Etienne, Ludovic Chahen, Pierre Louis Carrette, José Dugay, Jérôme Vial

► To cite this version:

Mohammad Shoeb, Girel Etienne, Ludovic Chahen, Pierre Louis Carrette, José Dugay, et al. New insight on the impact of NOx on diethanolamine (DEA) degradation mechanisms. International Conference on Greenhouse Gas Control Technologies, GHGT-17, Oct 2024, Calgary, Canada. 10.2139/ssrn.5069645. [hal-05036975](https://hal-05036975)

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17th International Conference on Greenhouse Gas Control Technologies, GHGT-17

20th -24th October 2024 Calgary, Canada

# New insight on the impact of NOx on diethanolamine (DEA) degradation mechanisms.

M. Shoeb<sup>a*</sup>, E. Girel<sup>a</sup>, L. Chahen<sup>a</sup>, P.L. Carrette<sup>a</sup>, José Dugay<sup>b</sup>, J. Vial<sup>b</sup>

<sup>a</sup>IFP Energies nouvelles, Rond-point de l'échangeur de Solaize, BP 3, 69360 Solaize, France

<sup>b</sup>Department of Analytical, Bioanalytical Sciences and Miniaturization (LSABM), UMR CBI 8231, ESPCI Paris, PSL Research University, 10 rue Vauquelin, 75005 Paris, France

## Abstract

Amine degradation presents a significant challenge in post-combustion CO₂ capture processes. While oxidative and thermal degradation pathways have been extensively studied, the impact of NOx-induced degradation remains relatively underexplored. Nitrosamines and nitramines are the primary degradation products formed through the reaction of amines with NOx; however, the mechanisms underlying their formation are not well understood, with various pathways proposed in the literature. In this study, the experimental degradation of diethanolamine (DEA) in the presence of NOx species (NO, NO₂, and their mixtures) was systematically investigated using specialized reactors. The effects of temperature on nitrosamine and nitramine formation were investigated. The degraded solutions were analyzed using Gas chromatography (GC-MS/FID) and ion chromatography (IC), coupled with a comprehensive mass balance assessment. The results obtained did not align with the proposed mechanism for nitrosamine formation by NO₂ gas. The stoichiometric ratio between nitrosamine and nitrate salt was found to be 3.6, rather than the expected 1. Our results suggested the possibility of alternative pathways for nitrosamine formation by NO₂ gas. No impact of temperature on nitrosamine formation was observed, except at 130 °C, where a high amount of nitrosamine was formed. In contrast, nitramine formation was directly influenced by temperature, with no nitramine formed at 130 °C. This highlights the need for further investigation to refine the understanding of NOx-induced degradation pathways and their impact on amine-based CO₂ capture.

**Keywords:** CO₂ post-combustion capture; amine degradation; diethanolamine; DEA; NOx impact; mechanisms, nitrosamines, nitramines.

## 1. Introduction

The continuous rise in carbon dioxide (CO₂) emissions has become a major driver of global climate change, prompting urgent calls for solutions to mitigate its environmental impact [1]. Post-combustion CO₂ capture (PCCC) technology, particularly with amine solvents, has emerged as a prominent and effective solution for capturing CO₂ from industrial emissions. This technology is widely adopted due to its proven efficiency and versatility in various industrial applications. However, a significant challenge with amine-based CO₂ capture systems is the degradation of the amine solvents over time, which compromises both the efficiency of CO₂ capture and increases operational costs [4,5]. In addition, the degradation products formed can be corrosive, carcinogenic, and pose environmental hazards [6].
While most of the research has focused on oxidative and thermal degradation processes, the impact of nitrogen oxides (NOx) in the flue gases on amine degradation remains relatively underexplored. NOx removal technologies, often implemented before CO₂ capture, typically achieve only up to 90% effectiveness [7]. Moreover, these systems require significant additional infrastructure and incur extra operational costs. Unlike CO₂, NOx does not readily release during CO₂ stripping operations and is irreversibly absorbed by the amine solvent. This leads to chemical reactions between NOx and amine solvents, resulting in the formation of nitrosamines and nitramines both of which are recognized carcinogens [6], as well as nitrite and nitrate salts as byproducts. These degradation products may not only pose health and environmental risks but also reduce the overall efficiency and effectiveness of the CO₂ capture process. This highlights a critical need to better understand the mechanisms of NOx-induced amine degradation, to develop more robust capture systems that mitigate these risks and improve the sustainability of post-combustion CO₂ capture technologies.

Several mechanisms have been proposed in the literature to explain the formation of nitrosamines and nitramines from the reaction of amines with NOx [2,3,8]. Nitrosamines are typically formed when an amine reacts with nitrite under acidic conditions (pH < 5). In 1957, White and Feldman demonstrated that nitrosamines, along with nitramines, could be readily formed from NOx gas in organic solvents [9]. However, their formation in aqueous media was not initially considered, as NOx gases were expected to hydrolyse in water rather than react directly with the amine. This view changed in 1976 when Challis and Kryptopoulos reported that nitrosamines and nitramines could indeed form in alkaline aqueous conditions, and that the reaction of NOx with amines is over 1000 times faster than its reaction with water [2,10,11].

NOx in flue gas primarily consists of NO and NO₂, typically in a 9:1 ratio, depending on the source [7]. According to the mechanism proposed by Challis and Kryptopoulos [2], as shown in Figure 1, NOx is initially absorbed into water (Figure 1, reaction I), leading to the formation of two reactive nitrogen species: dinitrogen trioxide (N₂O₃) and dinitrogen tetroxide (N₂O₄). N₂O₃ exists in two isomeric forms: symmetric (sym-N₂O₃) and asymmetric (asym-N₂O₃), both of which can act as nitrosating agents. The less stable isomer sym-N₂O₃ reacts nucleophilically with amine to form nitrosamine and nitrite as the byproduct (Figure 1, Reaction IIa). While the reaction can also proceed via asym-N₂O₃, this is not the predominant pathway. Similarly, N₂O₄ has symmetric (sym-N₂O₄) and asymmetric (asym-N₂O₄) isomers, which can act as both nitrosating and nitrating agents, depending on the isomer. The less stable isomer asym-N₂O₄ reacts with amine to form nitrosamine and nitrate as the byproducts (Figure 1, Reaction IIb), while the more stable isomer sym-N₂O₄ follows a different reaction pathway, leading to the formation of nitramine through a four-centred concerted mechanism, with nitrite as the byproduct (Figure 1, reaction IIc). In each case, for every mole of N₂O₄ gas that reacts with the amine, the reaction yields either nitrosamine and nitrate or nitramine and nitrite salts as byproducts, in a 1:1 molar ratio.

In addition, NOx reactions with amine solvent in carbon capture and storage (CCS) process have been investigated. Two mechanisms have been proposed: Fine *et al.* [3] proposed a NO₂-driven radical mechanism for amine nitrosation involving NO and NO₂ mixture (Figure 1, reaction IId). The initial steps involve the free radical absorption of NO₂ as nitrite, driven by abstracting a proton from the amine to form a free amine radical. This amine radical can subsequently react with NO and NO₂, leading to the formation of nitrosamine and nitramine. Ning *et al.* [8] proposed a molecular mechanism based on N₂O₃, similar to that proposed by Challis and Kryptopoulos (Figure 1, reaction IIa), as discussed earlier [2].
I) Formation of N₂O₃ and N₂O₄ in aqueous phase.

$$a) \text{NO(g)} + \text{NO}_2\text{(g)} \rightleftharpoons \text{N}_2\text{O}_3\text{(l)}$$

$$b) 2\text{NO}_2\text{(g)} \rightleftharpoons \text{N}_2\text{O}_4\text{(l)}$$

II) Nitrosamine and nitramine formation mechanism.

a) Nitrosation via sym-N₂O₃.

$$\underset{\substack{\text{R}' \\ \text{R}''}}{\text{NH}} + \text{N}_2\text{O}_3 \longrightarrow \left[ \underset{\substack{\text{R}' \\ \text{R}''}}{\text{NH}} \xrightarrow{\text{ON}-\text{ONO}} \text{R}' \text{N}-\text{NO} \right]^\ddagger \longrightarrow \underset{\substack{\text{R}' \\ \text{R}''}}{\text{NH}} + \text{NO}_2^- + \text{H}^+$$

b) Nitrosation via asym-N₂O₄.

$$\underset{\substack{\text{R}' \\ \text{R}''}}{\text{NH}} + \text{N}_2\text{O}_4 \longrightarrow \left[ \underset{\substack{\text{R}' \\ \text{R}''}}{\text{NH}} \xrightarrow{\text{ON}-\text{ONO}_2} \text{R}' \text{N}-\text{NO} \right]^\ddagger \longrightarrow \underset{\substack{\text{R}' \\ \text{R}''}}{\text{NH}} + \text{NO}_3^- + \text{H}^+$$

c) Nitration via sym-N₂O₄.

$$\underset{\substack{\text{R}' \\ \text{R}''}}{\text{NH}} + \text{N}_2\text{O}_4 \longrightarrow \left[ \underset{\substack{\text{R}' \\ \text{R}''}}{\text{O}_2\text{N}-\text{NO}_2} \right]^\ddagger \longrightarrow \underset{\substack{\text{R}' \\ \text{R}''}}{\text{N}-\text{NO}_2} + \text{NO}_2^- + \text{H}^+$$

d) NO₂-driven radical mechanism for nitrosation and nitration.

$$\underset{\substack{\text{R}' \\ \text{R}''}}{\text{NH}} + \text{NO}_2 \xrightarrow{\text{HNO}_2} \underset{\substack{\text{R}' \\ \text{R}''}}{\text{N}} + \underset{\substack{\text{R}' \\ \text{R}''}}{\text{NO}} + \underset{\substack{\text{R}' \\ \text{R}''}}{\text{N}-\text{NO}_2}$$

<caption>Figure 1: Proposed mechanisms of nitrosamine and nitramine formation by NOx [2,3].</caption>

In this study, we investigated the impact of NOx gases on diethanolamine (DEA) under various conditions. DEA was selected as a modelled secondary amine because, like any other secondary amines, it can react with NOx to form stable secondary nitrosamines. While primary and tertiary nitrosamines are unstable and prone to degradation, secondary nitrosamines tend to be more stable. Influence of temperature from 50 °C to 130 °C on nitrosation and nitration have also been investigated.
## 2. Experimental

### 2.1. Materials

Diethanoamine 99% (Thermo Fisher Scientific), N-nitrosodiethanolamine 97%+ (TCI Europe), NO₂ 2000 ppm diluted in N₂ (Air products), NO 2000 ppm diluted in N₂ (Air products). NO₂ 200 ppm and NO 1800 ppm mixture diluted in N₂ (Air products). Sodium nitrite 98% (Alfa Aesar), Sodium nitrate 99%+ (VWR chemicals), 69.5% HNO₃ (Carlo Erba reagents). Materials were used directly without any further purification.

### 2.2. Method

<figure><img src="image_3.png" /><figcaption>Figure 2: Reactor setup with labelled components.</figcaption></figure>

In this study, specialized reactors were used to investigate the degradation of DEA in the presence of NOx gas. The system consisted of six identical individual reactors, each designed to simulate the conditions of a CO₂ capture plant. The labelled scheme of the single reactor is shown in Figure 2. These reactors could operate independently under absorber and desorber conditions or in a cyclic mode. Each reactor was constructed from Hastelloy, had a capacity of 150 mL, and was equipped with essential components, including a condenser, pressure regulator, flowmeters, agitator, heater, and sampling system. The reactors were designed to operate at temperatures up to 150 °C and pressures of up to 30 bar.
For each experiment, the reactor was filled with 100–120 g DEA aqueous solution. The solution was degassed inside the reactor by purging with nitrogen gas (N₂) for approximately 30 minutes to remove dissolved CO₂ and oxygen and prevent oxidative degradation. After degassing, the solution was heated to the desired temperature. The reactor operated as an open system with continuous absorption, where NOx was injected continuously without a stripping step. To monitor the reaction, samples were taken regularly. At the end of each test, the remaining solution in the reactor was collected and weighed to determine any mass loss. Throughout the experiment, the pressure was maintained by the pressure regulator, while a condenser at the top of the reactor enabled the reflux of evaporated amine solvent and water. Operating parameters, including temperature, pressure, agitation, and gas inlet flow rates, were controlled and continuously monitored via iHistorian software.

## 2.3. Mass balance

The nitrogen and oxygen mass balances were evaluated based on the total amount of NO₂ injected and the products formed exclusively from NO₂, such as nitrosamine, nitramine, nitrite, and nitrate. Volatile nitrogen- or oxygen-containing degraded compounds in the gas phase were not analysed and are therefore not considered here. The total amount of NO₂ injected was determined using the flow rate, NO₂ concentration, and the duration of the experiment. The nitrogen mass balance was calculated as the ratio of the sum of nitrogen moles in the products to the total amount of NO₂ injected. The oxygen mass balance was calculated by accounting for the oxygen contribution from NO₂ present in the products.

$$NO_2 \text{ injected} = \frac{\text{Flow rate} \times \text{Time} \times NO_2 \text{ conc}}{22.4 \text{ L}}$$

$$Nitrogen\% = \frac{Nitrosamine + Nitramine + Nitrite + Nitrate}{Total NO_2} \times 100$$

$$Oxygen\% = \frac{Nitrosamine + 2 * Nitramine + 2 * Nitrite + 3 * Nitrate}{2 * Total NO_2} \times 100$$

## 2.4. Analysis

In this study, gas chromatography (GC) and ion chromatography (IC) were used to analyse the degraded amine solutions, with only liquid samples analysed. Volatile degraded compounds in the gas phase were not analysed.

### 2.4.1. Gas chromatography

Gas chromatography (GC) was used to identify amines and their degradation products, such as nitrosamines and nitramines. Degradation products were analysed using a TRACE 1600 Gas Chromatograph coupled with an ISQ 7610 Single Quadrupole Mass Spectrometer (GC-MS), fitted with a Thermo Scientific AI1610 Autosampler. An Agilent HP-5MS column was used for the separation. Quantification of nitrosamines and nitramines was carried out using an Agilent 6890N Gas Chromatograph with Flame Ionization Detection (GC-FID), equipped with an Agilent 7683B Injector and an Agilent CPSIL8-CB-Amines column. The uncertainty of the quantification method was ±10%. Detailed information about the analytical methods and GC specifications can be found in appendix section A.1. in Table 1. All samples were quantified by using external calibration, without dilution. NDELA standard was used for quantification, and NO₂DELA was quantified using a calibration plot of NDELA.

### 2.4.2. Ionic Chromatography

Formate, nitrite and nitrate concentrations were quantified using a Dionex ICS-6000 Ion Chromatograph (IC) equipped with a conductivity detector. The precolumn and main column used were Dionex IonPac-AG15 (4 × 50 mm)
and AS15 (4 × 250 mm), respectively. Aqueous KOH, with concentrations ranging from 8 to 60 mM, was used as the eluent, with a flow rate of 1 mL/min. Samples were appropriately diluted in ultra-pure water to fit within the calibration range of 0.5–10 ppm. Quantification was carried out with an uncertainty of ±10%.

# 3. Results and discussion

## 3.1. Without NOx

The degradation of DEA in absence of NOx was first investigated to assess whether any degradation products would form when diethanolamine was heated to high temperatures. A 30% (w/w) DEA aqueous solution was heated at 100 °C for 9 days. After the 9 days, no degradation of DEA was observed, indicating that under these conditions, DEA did not form any degradation products at 100 °C. This result also suggests that the solution was adequately degassed and there were no air leaks in the reactor during the experiment. If any air leaks had occurred, oxidative degradation products would have been detected.

## 3.2. With NOx

The impact of NO, NO₂, and their mixture on the degradation of DEA was investigated using 30% (w/w) DEA aqueous solution at 70 °C, 10 bar pressure and 160 h, with different NOx concentrations of 2000 ppm NO diluted in N₂, 2000 ppm NO₂ diluted in N₂, 1800 ppm NO and 200 ppm NO₂ diluted in N₂, at a flow rate of 5 NL/h. The Figure 3 shows the GC-FID chromatograph of the degraded DEA under different NOx conditions.

<figure><img src="image_4.png" /><figcaption>Figure 3: GC chromatograph of degraded DEA under different NOx conditions. Experimental conditions: 30% DEA, 0.2% NOx diluted in N₂, 5 NL/h, 70 °C, 160 h, 10 bar.</figcaption></figure>

When only NO was used, no degradation of DEA was observed. The small peaks present in the chromatograph corresponded to impurities in the starting DEA and method artifacts, which were also detected in the blanks. The use of a mixture of NO and NO₂ resulted primarily in the formation of N-nitrosodiethanolamine (NDELA), consistent with the literature [3,8]. When both NO and NO₂ were present, with NO in a significantly higher concentration than NO₂, as in our case with a 9:1 NO to NO₂ ratio, the reaction predominantly led to the formation of nitrosamines. In contrast, when NO₂ alone was used, several degradation products were observed. These included monoethanolamine
(MEA), NDELA, and N,N-bis(2-hydroxyethyl)formamide (BHF), all of which were identified using GC-MS and the NIST library, and nitrodiethanolamine (NO₂DELA), which was confirmed by comparing the mass fragmentation pattern with the reported literature [12].

Nitrosamine and nitramine are expected degradation products from the NO₂-induced degradation of DEA. Other observed products are typically associated with oxidative degradation of amines, but since oxygen was not present in our system, it suggests that these products may have resulted from secondary reactions involving oxidation of amine by NO₂. Chih-I Chen *et al.* has reported and proposed a mechanism for the formation of degradation products, other than nitrosamines and nitramines, resulting from the oxidation of piperazine by NO₂ [13]. NO₂ can abstract protons from amines forming amine radicals. These amine radicals may then propagate and further oxidize the amines [3,13].

MEA is a well-known degradation product of DEA, as reported in the literature [4]. To the best of our knowledge, while BHF has not been previously reported in DEA degradation studies, a similar oxidative degradation product, N-(2-hydroxyethyl)formamide (HEF), has been reported in the degradation of MEA [4]. HEF is formed by the reaction of MEA with formic acid[14]. Formic acid is also a degradation product of DEA and MEA [4], was found at a concentration of approximately 500 ppm at the end of experiment. In our case, BHF may have formed through a similar mechanism to that of HEF, as illustrated in Figure 4.

<figure><img src="image_5.png" /><figcaption>Figure 4: Mechanism of N,N-bis(2-hydroxyethyl)formamide (BHF) formation.</figcaption></figure>

These results indicated that NO₂ has more adverse effect in amine degradation compared to NO and NO₂ with an excess of NO. DEA degradation with NO₂ was further investigated by analysing the products and by-products formed, considering reaction stoichiometry, mass balance, and comparisons with the literature.

<figure><img src="image_6.png" /><figcaption>Figure 5: A) Accumulation of degraded products over time. B) Mass balance at the end of reaction. Experimental conditions: 30% DEA, 2000 ppm NO₂ diluted in N₂, 5 NL/h, 70 °C, 10 bar.</figcaption></figure>
Nitrosamine, nitramine, nitrite, and nitrate are the primary products resulting from the reaction of NO₂ with DEA. These species were monitored and quantified as they accumulated in the solvent over time. Figure 5A illustrates the accumulation of NDELA, NO₂DELA, nitrite, and nitrate in mmol/kg over time, Figure 5B shows the nitrogen and oxygen mass balance in mole percentage at the end of reaction.

Nitrite was formed in the highest amount, representing 62% of the total NO₂ injected, followed by nitrosamine at 22%, nitramine at 15%, and nitrate at 6%. Together, these species accounted for all the NO₂ injected within a 10% uncertainty, thereby closing the nitrogen and oxygen mass balance. This confirms that all the NO₂ injected was reacted and converted into these species, with no significant other species directly originating from the NO₂ reaction being unaccounted for.

Our results do not align with the mechanism proposed by Challis and Kryptopoulos for nitrosation by NO₂. If we consider free radical absorption of NO₂ as nitrite and forming amine radicals (Figure 1 reaction IId) (8), the expected product would be only nitramine, resulting from the reaction of the amine radical with NO₂. The formation of nitrosamine would be unlikely, as NO is absent in the system. Additionally, nitrite typically reacts with amine carbamate to form nitrosamine under high temperatures in stripper conditions [15,16]. Nitrosamine formation at high temperature with unloaded amine have also been reported [17]. In our case, DEA was not loaded with CO₂, and the reaction temperature of 70 °C may not have been sufficient to drive this reaction. However, the higher formation of nitrosamine compared to nitramine suggests an alternative mechanism, potentially involving the formation of N₂O₄, as proposed by Challis as discussed earlier [2]. According to the proposed mechanism described in (Figure 1, reaction IId), for each mole of N₂O₄ reacting with amine, one mole of nitrosamine and one mole of nitrate should be formed, suggesting a 1:1 ratio between nitrosamine and nitrate. However, our results showed a higher amount of nitrosamine formed compared to nitrate, with a molar ratio of 3.6. Repeated experiments consistently yielded ratio higher than one. One possible hypothesis was that nitrate could have been consumed during the reaction to form additional degradation products. To investigate this hypothesis, a 30% (w/w) DEA aqueous solution containing 50 mmol each of NaNO₂ and NaNO₃ was heated at 100 °C in a closed reactor for 7 days. No degradation products or nitrosamine formation were observed. The similar experiment was conducted using only NaNO₂ at a much higher temperature of 150 °C, yielding the same result, i.e. no nitrosamine formation. This suggests that neither nitrite nor nitrate could react with the amine under the conditions tested, which does not explain the low levels of nitrate observed. These results suggest that the mechanism of NO₂ reaction with amine leading to the formation of nitrosamine may be different from what was previously proposed, or it might involve more than one mechanism in addition to the N₂O₄ mechanism.

The observed ratio of nitramine to nitrite was not 1:1, which could be attributed to the free radical absorption of NO₂ as nitrite (Figure 1, IId) [3], in addition to the nitration reaction where nitrite is also formed as a byproduct. Some part of the injected NO₂ may have been absorbed as free radical, forming amine radicals and nitrite. This could explain the high amount of nitrite present, as well as the formation of other degradation products observed such as MEA and BHF. The free amine radicals formed might have reacted with each other, degraded, or interacted with other species, leading to the formation of those products [13].

### 3.3. *Impact of temperature.*

The influence of temperature on the formation of nitrosamines and nitramines was investigated within a temperature range of 50 °C to 130 °C. These experiments were conducted under modified conditions compared to previous studies, using 10.5% (w/w) DEA aqueous solution, 2000 ppm NO₂ diluted in N₂, flow rate of 20 NL/h, duration 30 h and a pressure of 3 bar. Figure 6 illustrates the effect of temperature on the formation of NDELA and NO₂DELA over time.
<figure><img src="image_7.png" /><figcaption>Figure 6: Effect of temperature (50 °C to 130 °C) on the formation of N-nitrosodiethanolamine (NDELA) and nitrodiethanolamine (NO₂DELA) over time. Experimental conditions: 10.5% (w/w) DEA, 2000 ppm NO₂ diluted in N₂, 20 NL/h, 30 h, 3 bar.</figcaption></figure>

At 130 °C, the formation of NDELA was the highest, while no clear or consistent relationship was observed at lower temperatures. The elevated NDELA levels at 130 °C may be attributed to the conversion of nitrite to nitrosamine at this temperature. Chandan *et al.* also did not observe any effect of temperature on nitrosomorpholine formation within the 20 to 40 °C range, attributing this to the balance between gas solubility in the solution and reaction kinetics [17]. This aligns with our results and may explain the similar amount of NDELA formed at temperatures up to 100 °C. Conversely, NO₂DELA formation decreased with rise in temperatures and was not formed at 130 °C. This might be attributed to the instability and degradation of NO₂DELA at higher temperatures or decrease in N₂O₄ formation as the temperature increases. Since the equilibrium between NO₂ and N₂O₄ in the gas phase is temperature-dependent, higher temperatures tend to shift the equilibrium toward NO₂ [18].

To check the stability of NO₂DELA, a solvent previously degraded with NO₂ was heated at 100 °C and 130 °C in the absence of NO₂, in an open reactor under N₂ gas bubbling for 65 hours. The concentration of NO₂DELA remained unchanged as shown in Figure 7, indicating its stability and resistance to degradation. In contrast, the concentration of NDELA increased slightly at 100 °C, and then rose rapidly at 130 °C. This increase was due to the reaction of nitrite with amine to form nitrosamine at high temperature. This indicates a strong temperature dependence of nitrosamine formation by nitrite, consistent with findings reported in the literature [17]. These findings suggest that NO₂ reacts differently with amine at elevated temperatures, resulting in less nitramine formation. Furthermore, the results imply that nitramine formed during the CO₂ absorption cycle may not degrade in the desorber column, even under high-temperature conditions.
<figure><img src="image_8.png" /><figcaption>Figure 7: Evolution of N-nitrosodiethanolamine (NDELA) and nitrodiethanolamine (NO₂DELA) over time. Experimental Conditions: 60 g degraded DEA solution, 100 °C and 130 °C, N₂ bubbling, 20 NL/h, 65 h, 3 bar.</figcaption></figure>

**4. Conclusion**

We have investigated the effect of NOx on DEA degradation under different conditions and found that NO₂ have the most adverse effect on amine degradation. DEA degradation by NO₂ resulted in the formation of nitrosamines, nitramines, and other byproducts. The discrepancies between the proposed mechanism in the literature for nitrosamine formation via NO₂ and our experimental results suggest the possibility of additional reaction pathways. Further research is underway with different amines to investigate the mechanisms.
# Appendix A.

## A.1. GC-FID and GC-MS analysis method and specification

Table 1 provides the detailed information about the GC methods.

<table><caption>Table 1: GC-FID and GC-MS conditions.</caption><thead><tr><th></th><th>GC/MS (EI)</th><th>GC-FID</th></tr></thead><tbody><tr><td>column</td><td>HP-5MS</td><td>CPSIL8-CB-Amines</td></tr><tr><td>length</td><td>30 m</td><td>25 m</td></tr><tr><td>internal diameter</td><td>250 mm</td><td>320 µm</td></tr><tr><td>thickness</td><td>0.25 µm</td><td>1.2 µm</td></tr><tr><td>Injection volume</td><td>0.2 µL</td><td>1 µL</td></tr><tr><td>Injection mode</td><td>Split</td><td>Split</td></tr><tr><td>Split ratio</td><td>45</td><td>30</td></tr><tr><td>Injector temperature</td><td>280 °C</td><td>280 °C</td></tr><tr><td>Oven initial temp</td><td>40</td><td>40</td></tr><tr><td>Initial hold time</td><td>0</td><td>0</td></tr><tr><td>Oven ramp</td><td>2 °C/min</td><td>2 °C/min</td></tr><tr><td>Oven final temp</td><td>200 °C</td><td>200 °C</td></tr><tr><td>final hold time</td><td>2</td><td>2</td></tr><tr><td>Carrier</td><td>He</td><td>He</td></tr><tr><td>flow rate (constant)</td><td>1.0 ml/min</td><td>1.6 ml/min</td></tr><tr><td>detector (FID) temp</td><td>-</td><td>280</td></tr></tbody></table>

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