Nitrosamines and Nitramines in Amine-Based Carbon Dioxide

Critical Review

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Nitrosamines and Nitramines in Amine-Based Carbon Dioxide Capture Systems: Fundamentals, Engineering Implications, and Knowledge Gaps Kun Yu, William A. Mitch, and Ning Dai Environ. Sci. Technol., Just Accepted Manuscript • Publication Date (Web): 25 Sep 2017 Downloaded from http://pubs.acs.org on September 25, 2017

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Environmental Science & Technology

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Nitrosamines and Nitramines in Amine-Based Carbon Dioxide Capture Systems:

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Fundamentals, Engineering Implications, and Knowledge Gaps

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Kun Yu1, William A. Mitch2, and Ning Dai3*

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Department of Chemical and Biological Engineering

University at Buffalo, The State University of New York, Buffalo, NY, 14260

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Department of Civil and Environmental Engineering Stanford University, Stanford, California 94305

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Department of Civil, Structural and Environmental Engineering

University at Buffalo, The State University of New York, Buffalo, NY, 14260

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*

Corresponding author: Phone: (716) 645-4015; Fax: (716) 645-3667 Email: [email protected]

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ACS Paragon Plus Environment


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Abstract

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Amine-based absorption is the primary contender for post-combustion CO2 capture from fossil

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fuel-fired power plants. However, significant concerns have arisen regarding the formation and

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emission of toxic nitrosamine and nitramine byproducts from amine-based systems. This paper

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reviews the current knowledge regarding these byproducts in CO2 capture systems. In the

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absorber, flue gas NOx drives nitrosamine and nitramine formation after its dissolution into the

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amine solvent. The reaction mechanisms are reviewed based on CO2 capture literature as well as

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biological and atmospheric chemistry studies. In the desorber, nitrosamines are formed under

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high temperatures by amines reacting with nitrite (a hydrolysis product of NOx), but they can

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also thermally decompose following pseudo-first order kinetics. The effects of amine structure,

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primarily amine order, on nitrosamine formation and the corresponding mechanisms are

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discussed. Washwater units, although intended to control emissions from the absorber, can

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contribute to additional nitrosamine formation when accumulated amines react with residual

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NOx. Nitramines are much less studied than nitrosamines in CO2 capture systems. Mitigation

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strategies based on the reaction mechanisms in each unit of the CO2 capture systems are

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reviewed. Lastly, we highlight research needs in clarifying reaction mechanisms, developing

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analytical methods for both liquid and gas phases, and integrating different units to quantitatively

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predict the accumulation and emission of nitrosamines and nitramines.

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INTRODUCTION

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Carbon capture and storage, the separation of CO2 from industrial emission sources and

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the subsequent geological storage, is a critical endeavor to mitigate climate change.1 Fossil fuel-

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fired power plants are the largest point sources of anthropogenic CO2 emissions,2 and will

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largely rely on post-combustion CO2 capture to reduce emissions.1,

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technology is the primary contender for post-combustion CO2 capture due to its long-term

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operational experience in pre-combustion CO2 separation from natural gas and hydrogen and its

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suitability for retrofitting existing power plants.4 Amine technology has been demonstrated at

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full scale, such as TMC Mongstad in Norway (300,000 tonnes per year CO2 captured) and BD3

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SaskPower in Canada (1 million tonnes per year CO2 captured).5-8

3

Amine-based absorption

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Amine-based CO2 capture systems are comprised of an absorber and a desorber (Figure

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1). In the absorber, flue gas CO2 is absorbed by an aqueous amine solution (i.e., “solvent”),

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which counter-currently scrubs the flue gas at 40–60 °C and achieves 75–90% CO2 removal.9

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The absorption is driven by the reactions between amines and CO2, forming carbamate in

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primary and secondary amine solvents (Table1, reaction 1) or bicarbonate in tertiary amine

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solvents (reaction 2). The CO2-rich solvent is then pumped to the desorber, within which the

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high temperature (typically 100–140 ºC)10 reverses reaction 1 or 2 to release CO2. The CO2-lean

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solvent is passed through a heat exchanger to pre-heat the CO2-rich solvent, and then circulated

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back to the absorber. Washwater units (also referred to as “water wash”) are single- or multi-

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stage water scrubbers often included in the upper part of the absorber to recover amines and

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control emissions of amine degradation products. Active research topics in amine-based CO2

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capture include improving energy efficiency, reducing amine loss, and minimizing

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environmental impacts.5, 11



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The formation and subsequent release of potentially carcinogenic nitrosamine and

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nitramine byproducts is a major concern regarding the environmental impacts of the amine

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systems.12-15 Nitrosamines and nitramines form in the absorber from reactions between NOx in

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the flue gas and amines in the solvent. Nitrosamines also form in the desorber from reactions

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between aqueous nitrite, a hydrolysis product of NOx, and amines.16-18 Meanwhile, nitrosamines,

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and with less direct evidence, nitramines, undergo decomposition in the desorber.19-21 Some

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nitrosamines can increase cancer risks at low nanogram per liter levels in drinking water.15

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Nitrosamines and nitramines have been detected in pilot-scale systems. In the CO2-lean

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solvent of a facility operated with monoethanolamine (MEA) at a capacity of 800 US tons

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CO2/day (Trona, CA), a total nitrosamine concentration of 2.91 mM was measured.22, 23 The total

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nitrosamine concentration in the 30% MEA solvent from a pilot system (Trondheim, Norway)

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was 4.8–12.3 µg/mL (approximately 48–123 µM).21 Nitrosamines and nitramines formed in the

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solvent can be stripped into the gas phase and enter the atmosphere via stack emissions, a major

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route for these byproducts to impact the environment. The exhaust from a pilot-scale MEA-based

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CO2 capture system (Rotterdam, Netherlands; maximum 250 kg CO2 captured per hour)

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contained 5–47 ng/Nm3 nitrosamines and nitramines.24 The washwater unit captures some of

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these emissions. Washwater samples from a pilot system without a desorber operated with a 25%

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2-amino-2-methyl-1-propanol/15% piperazine solvent and a 35% MEA solvent contained 59 µM

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and 0.73 µM total nitrosamines, respectively.25 An example of potential regulatory standards was

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the discharge permit issued by the Norwegian Climate and Pollution Agency for the proposed

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full-scale facility at CO2 Technology Centre Mongstad, which limited the sum of all

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nitrosamines and nitramines to 0.3 ng/m3 in downwind airsheds and 4 ng/L in downwind water

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supplies.26


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The goal of this work is to critically review the current knowledge on the formation and

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decomposition of nitrosamines and nitramines in amine-based CO2 capture systems. In addition

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to recent research on CO2 capture systems, previous biological and atmospheric chemistry

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studies on nitrosamine and nitramine formation mechanisms are critically examined for their

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applicability to amine-based systems, although they frequently employed distinct reaction

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conditions from those of CO2 capture systems (Table 2). Nitrosamines and nitramines may also

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form in the atmosphere when the amines emitted by the CO2 capture systems react with ambient

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NOx. This topic has been reviewed by Nielsen et al.13 and Ge et al.27, and is not covered here.

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In this review, we first discuss the NOx-driven formation of nitrosamines and nitramines

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in the absorber, with special attention paid to the relative importance of NO and NO2. The

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involvement of other flue gas constituents, such as CO2 and SO2, is also considered. Second, we

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evaluate the desorber regarding the formation of nitrosamines from nitrite and their

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decomposition. Third, we discuss the influence of solvent amine structure on the formation of

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nitrosamines and nitramines and the distribution of specific nitrosamine products. Fourth, we

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review the importance of washwater units for capturing volatile contaminants in the absorber

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exhaust, and circumstances in which washwater units can serve as additional sources of

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nitrosamines. Based on the formation and decomposition mechanisms in each unit, a few

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mitigation options for nitrosamines and nitramines are reviewed. Lastly, we present a research

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outlook identifying major gaps in the field.

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ABSORBER: Heterogeneous NOx–Amine Reactions to Form Nitrosamines and

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Nitramines



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The composition of power plant flue gases varies with fuel source and plant operational

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conditions. Flue gases from coal-fired power plants typically contain 14% CO2, 5% O2, 81% N2,

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300–3000 ppmv SOx, and 100–1000 ppmv NOx, while those from natural gas-fired power plants

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typically contain 4% CO2, 15% O2, 81% N2, less than 1 ppmv SOx, and 100–500 ppmv NOx.28

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Flue gas NOx is primarily NO (90–95%), with approximately 5–10% NO2.29 NOx and SOx

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control units are typically implemented prior to CO2 capture. NOx control technologies can

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achieve 90% NOx reduction, but NOx composition after treatment depends on the control

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technologies.30-32

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2.1

115

Scale Systems

NOx Dependence of Nitrosamine and Nitramine Formation in Laboratory- and Pilot-

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In laboratory-scale absorbers using morpholine (MOR) or piperazine solvents treating

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synthetic flue gas (6–7.5 L/min, 25–150 ppmv NOx), nitrosamine accumulation rates were 1.3 ×

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10-5–3.6 × 10-4 g/min (1.1 × 10-7–3.1 × 10-6 mol/min).33,

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determine the relative contribution of the two NOx species, NO and NO2, to nitrosamine

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formation. Under absorber conditions, when only NO or NO2 is present, NO2 is more potent than

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NO for forming nitrosamines;33, 35, 36 however, when NO and NO2 coexist, nitrosamine formation

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exhibits a first-order dependence on both.33,

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continuously purged through a 5 M MOR solution for 6 hours at 30 ºC, 88 µg/g

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N-nitrosomorpholine (NMOR) formed; only 12–22 µg/g of NMOR formed when 500 ppmv NO

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was used.35, 36 Similarly, five times faster NMOR accumulation was observed with 3 ppmv NO2

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than with 15 ppmv NO in a packed-bed column reactor.33 However, in the same column reactor,

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the NMOR accumulation rate linearly increased with NO concentration (5–22.5 ppmv) with a

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fixed NO2 concentration of 3 ppmv, as well as with NO2 concentration (3–15 ppmv) with a fixed

34

34

Researchers have attempted to

In a batch system, where 100 ppmv NO2 was


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NO concentration of 15 ppmv.33 In the presence of 3 ppmv NO2, the NMOR accumulation rate

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was 8 times higher with 22.5 ppmv NO than without NO.33 Because NO and NO2 coexist in

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authentic flue gases, the latter observations are more relevant to industrial absorbers. A first-

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order dependence on gaseous NO2 concentration was also reported for 1-nitrosopiperazine

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formation in piperazine solvents, although the dependence on NO was not evaluated.34

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Two pathways have been proposed for nitrosamine formation in the absorber. Fine and

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Rochelle proposed an NO2-initiated radical pathway, in which an amine radical is first formed by

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NO2 abstracting the amino-hydrogen atom, and the amine radical subsequently reacts with NO to

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form a nitrosamine (Scheme 1a).34 Two pieces of evidence were used to support this pathway: 1)

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the detection of 2-piperazinol, a transformation product of the amine radical, in piperazine

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solvents, and 2) the inhibition of 1-nitrosopiperazine formation by a proprietary radical

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scavenger Inhibitor A.34 In contrast, Dai and Mitch33 proposed that N2O3, the reactive

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intermediate formed by NO and NO2 (reaction 3), was the main species contributing to

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nitrosamine formation in the absorber; two pieces of evidence were put forth. First, the NMOR

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formation rate in MOR solvents has a first-order dependence on both NO and NO2

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concentrations, suggesting that the rate-limiting step involves both species. Second, the

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formation of nitrosamines far exceeds nitramines,33 which is contradictory to that anticipated

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from the NO2-initiated radical pathway. Once an amine radical is formed, it is expected to

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reactive non-selectively with NO and NO2, which would lead to similar formation of

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nitrosamines and nitramines given that the estimated aqueous concentrations of NO and NO2

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were similar.33, 34 Each of these two studies has their limitations. In the former, the influence of

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NO concentration on nitrosamine formation was not evaluated, nor the possibility of the

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proprietary radical scavenger Inhibitor A hindering nitrosamine formation by scavenging NO



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and NO2, both radicals. In the latter, gas phase NOx concentrations were not quantified,

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rendering it difficult to establish reaction stoichiometry or mass balance. Further discussion on

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the NO2-initiated radical pathway and the N2O3 pathway is included in section 2.2, with

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reference to biological and atmospheric chemistry studies.

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Because NO2 can be generated via oxidation of NO by O2 (reaction 4), the contribution of

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this reaction to nitrosamine and nitramine formation was considered. In the absence of NO2, 1-

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nitrosopiperazine and N,N’-dinitrosopiperazine formed in a piperazine solution purged with 8000

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ppmv NO and 5% O2,37 suggesting that NO oxidation can lead to nitrosamine formation.

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However, this contribution may be limited to scenarios with high concentrations of NO. With a

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fixed 9:1 NO to NO2 ratio, Fostås et al. observed that

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N-nitrosodiethanolamine (NDELA) in MEA solvent increased with O2 (0−14%) when the total

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NOx concentration was 25 ppmv or 45 ppmv, but not at 5 ppmv.17 Similarly, Dai and Mitch did

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not observe a dependence on oxygen for NMOR accumulation in the absorber at low NOx

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concentrations (15 ppmv NO and 3 ppmv NO2).33 While these results suggest nitrosation via NO

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oxidation is a minor pathway, the formation of NMOR was detectable in the presence of NO

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alone,33, 35, 36 which is consistent with a computational study.38 Thus, pre-treatment of flue gases

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to selectively remove NO2 may reduce, but not completely eliminate nitrosamine formation.

the accumulation of

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The relative accumulation of nitrosamines and nitramines is dependent on NOx

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concentration and composition, but nitramines are generally the minor product in the solvents

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tested under conditions relevant to flue gases (5–10% of total NOx being NO2).17,

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Nitramine formation increases with flue gas NO2 but not with NO.33 In a packed-bed column

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reactor, NMOR accumulation exceeded N-nitromorpholine (NO2-MOR) by approximately an

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order of magnitude with varying NO (0–22.5 ppmv) and NO2 (0–15 ppmv) concentrations.33 In a


25, 33, 34

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high gas flow apparatus with 100 ppmv NO and 10 ppmv NO2, Fine and Rochelle determined

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the yield of 1-nitrosopiperazine and 1-nitropiperazine from absorbed NOx to be 12% and under

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5% (below detection limit), respectively.34 In a pilot system operated with MEA solvent, the

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nitramine 2-(nitroamino)ethanol accumulated at levels approximately three orders of magnitude

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lower than total nitrosamines under 5 or 50 ppmv NOx (NO to NO2 ratio not reported).21 Even

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with 25 ppmv equimolar NO and NO2, NO2-MOR accumulation was ten times slower than

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NMOR.25 Only when NO2 is present alone does formation of nitramines become comparable

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with nitrosamines. When 25 ppmv NO2 was purged through a 5 M MOR solution, the formation

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rate of NO2-MOR was one third that of NMOR.25 When 100 ppmv NO2 was purged through 0.5–

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2.5 M MOR solutions, NO2-MOR formed at comparable levels to NMOR with less than 20%

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difference.35 These conditions, however, are not relevant to typical flue gases.

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Lastly, it is worth mentioning that the accumulation of nitrite in the absorber is also

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dependent on NOx concentrations (reactions 8–10). Although nitrite does not form nitrosamines

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in the absorber, it becomes a nitrosating agent in the desorber (as further discussed in section

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3.1).35 Nitrite accounted for the majority of NOx absorbed in the absorber,34 and accumulated

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more than ten times faster than nitrosamines across different amine solvents, gas compositions,

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and absorber configurations.16,

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NO2,33, 34 but not with NO.33 Nitrate, the other NOx hydrolysis product, forms more slowly than

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nitrite in the absorber under typical flue gas NOx compositions (high NO to NO2 ratio),34 but

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accumulates at higher concentrations due to its stability in the desorber.22, 39 Nitrate is not known

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to contribute to nitrosamine or nitramine formation in CO2 capture systems.

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2.2

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and HONO

33, 34, 36

Nitrite accumulation linearly increased with flue gas

Nitrosamine and Nitramine Formation Mechanisms: The Role of NO2, N2O3, N2O4,



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Prior to the interest in nitrosamine and nitramine formation during CO2 capture,

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biological and atmospheric chemistry research also studied NOx-amine interactions for the

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formation of nitrosamines and nitramines. These studies concluded that N2O3 is a nitrosating

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agent, while NO2 and N2O4 can initiate both nitrosation and nitration reactions.40-48 HONO was

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considered a nitrosating agent in the atmosphere.49,

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reviewed, noting their strengths and limitations by considering the similarities and differences of

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biological (physiological) and atmospheric systems from CO2 capture systems. As shown in

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Table 2, amine solvents feature basic pH: the CO2-lean and CO2-rich (0.2 C/N) MEA solvents

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are at pH 13 and pH 10, respectively,51 while gastric and physiological pH values are 3–4 and

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7.4, respectively.44, 45 Amine concentrations in the solvents are 2.5–8 M (i.e., 15–50 wt%),9 far

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exceeding those in biological systems (10–3–10–2 M)39-41 and in the atmosphere (< 10 ppbv),52, 53

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which may allow different reaction pathways. In addition, amine systems operate at elevated

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temperatures (e.g., absorber 40–50 ºC and desorber 120 ºC) and are designed to facilitate mass

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transfer between the gas and liquid phases.

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This body of literature is critically

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NO2 vs. N2O3 as nitrosating agents: Mirroring the debates in the CO2 capture literature,

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the relative importance of NO2 and N2O3 as gaseous nitrosating agents has been a contentious

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topic.41-44,

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earlier studies in developing the corresponding reaction mechanisms, a chronological approach is

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taken here to trace the evolution of theories in order to facilitate a critical review on their

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applicability to CO2 capture systems. Nitrosation was most extensively studied in aqueous acidic

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nitrite solutions simulating the gastric environment, in which N2O3, in equilibrium with nitrous

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acidium ion (H2ONO+) (reaction 5), was considered the active nitrosating agent.45, 59, 60 However,

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these studies do not directly apply to the absorber where nitrosating agents originate from the gas

47, 54-58

Because current CO2 capture literature tends to reference a subset of these


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phase. Therefore, studies employing gaseous NOx to react with amines in aqueous solutions are

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the focus of this section. Challis and colleagues were among the first to explore this topic.47, 48, 54,

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57

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heterogeneous N2O3 pathway was responsible (Schemes 1b and 1c).47, 55, 57 In their experiments,

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a series of amines were dissolved in 0.1 M NaOH aqueous solutions and exposed to NOx (1000–

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100,000 ppmv) in the headspace; the extent of nitrosation was monitored.47,

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amines of drastically different basicity (pKa 1–11) were nitrosated to similar extents,54

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contradictory to the positive correlation between basicity and reactivity expected from

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nucleophilic attack on N2O3,45 it was tentatively proposed that NO2 radical initiated the

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nitrosation reactions by forming an amine radical (Scheme 1a).54 However, further investigation

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by the same researchers showed that the amine radical is unlikely to be a major intermediate,

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because the decomposition products normally associated with amine radicals were absent.47

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Accordingly, they proposed the heterogeneous N2O3 pathway (Schemes 1b and 1c), which

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involves the reaction of ON-ONO, one of the isomers of N2O3, with amines in the aqueous

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phase.47, 55, 57 This isomer forms when NO and NO2 combine in the aqueous phase after their

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dissolution from the gas phase (Scheme 1b).47 In particular, by drawing an analogy to the known

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relative stability of the two N2O3 isomers in the gas phase, it was hypothesized that in the

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solution phase, the ON-ONO isomer of N2O3 is more reactive than its ON-NO2 isomer.47 The

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ON-NO2 isomer is the dominant form of N2O3 in acidic nitrite solutions; its relatively low

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reactivity could explain the absence of nitrosation for weakly basic amines (pKa ≤ 2) in acidic

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nitrite solutions, in contrast to facile nitrosation for these amines when exposed to gaseous

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NOx.57 The aqueous combination of NO and NO2 to form the ON-ONO isomer was considered

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the rate limiting step for nitrosation.47

They initially proposed the NO2-iniated radical pathway (Scheme 1a),54 but later stated that a


48, 54, 57

Because


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Later studies by Cooney et al. expanded the experimental pH range to near neutral and

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employed lower NOx concentrations. When mixtures of NO (0–90 ppmv) and NO2 (4–26 ppmv)

246

were purged through 5 mL of 10 mM MOR solution, first-order dependencies on both NO and

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NO2 were observed for NMOR formation.42, 43 While the authors argued that the NO2-initiated

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radical pathway predominated, the experimental results used to support this argument were from

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NO2 purging experiments in the absence of NO, which naturally excluded the contribution of

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N2O3.40,41

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The next group of studies considered coupled reactions of NO oxidation to NO2 and

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amine nitrosation; the findings reinforced the importance of N2O3 for nitrosation of amines.41, 44,

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46

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thiols (0.5–150 mM, pH 7.4) and one with dissolved O2, and monitored the accumulation of

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reaction products using a stopped-flow apparatus.41 They reported that the formation of NMOR

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or S-nitrosothiols was second-order with respect to NO and first-order with respect to O2,

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suggesting that NO autoxidation was the rate-limiting step. They also proposed that, following

258

NO oxidation to NO2, the NO2-initiated radical pathway and the N2O3 pathway were two parallel

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pathways leading to nitrosation. Using kinetic data, they showed that although some thiols were

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primarily nitrosated through the NO2-initiated radical pathway, other thiols and MOR were

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nitrosated through the N2O3 pathway. MOR, in particular, appeared to be at least 10 times less

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likely to form a radical by NO2 than most thiols.41 Lewis et al.44 and Caulfield et al.46 also

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concluded that N2O3 was the responsible nitrosating agent. They introduced gaseous NO and O2

264

via purging to a MOR solution (0.05−2.5 mM, pH 7.4) and observed inhibition on nitrosation by

265

phosphate and chloride ions, presumably through scavenging of N2O3 (reaction 6).44, 46 Although

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these two groups of studies differ in experimental setup in that the former pre-dissolved NO and

Goldstein and Czapski mixed two aqueous solutions, one with dissolved NO and MOR or


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O2 into aqueous solutions41 while the latter introduced the two gaseous species via gas purging,44,

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46

both reached agreement that N2O3 was responsible for MOR nitrosation.

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Recently, computational tools such as density functional theory (DFT) calculations were

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applied to investigate nitrosation by NO2 and N2O3. Most of these studies were limited to the gas

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phase, but limited evaluations considered reactions in the aqueous phase. For gas phase

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reactions, Zhao et al.61 and Sun et al.62 supported the NO2-initiated radical pathway and the N2O3

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pathway for gas phase reactions, respectively. Both studies suggested that electron-donating

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moieties incrased the reactivity of amines,61,

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Challis et al.57 Another study by Sun et al. compared the reactivity of the N2O3 isomers, sym and

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trans-cis N2O3 (both ON-ONO) and asym N2O3 (ON-NO2), in both the gas and aqueous phases,

277

and showed that the energy barrier for the former is close to or even lower than that of the

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latter.63 This is consistent with the mechanism proposed by Challis et al.47

62

in contrast to the experimental results from

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In summary, the relative importance of NO2 and N2O3 as nitrosating agents is a

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contentious topic in the biological and atmospheric chemistry literature, similar to that in the

281

CO2 capture literature, but there is more favorable evidence supporting N2O3, formed in the

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aqueous phase following dissolution of gas-phase NOx, as the major nitrosating agent of amines

283

within the aqueous phase. Further research is needed to verify the nitrosation reaction

284

mechanisms in the absorber of CO2 capture systems. Lastly, it is worth mentioning that N2O4,

285

while behaving similarly to N2O3 with respect to nitrosation,47 is less relevant for CO2 capture

286

systems, because its formation is not favored at the high NO/NO2 ratios and relatively low NO2

287

concentration typical of flue gases, and hence only its relevance to nitration is reviewed below.

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NO2 vs. N2O4 as nitrating agents: DFT calculations for dimethylamine showed that

289

nitration reactions in the aqueous phase may proceed through either NO2 or N2O4,64 but



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experiments using morpholine, pyrrolidine, or N-methylaniline suggest that the primary

291

mechanism is the NO2-initiated radical pathway (Scheme 1a)42, 48. Regarding N2O4, its O2N-NO2

292

isomer can be subject to nucleophilic attack by an amine to form a nitramine (Scheme 1d). In 2

293

mM MOR or pyrrolidine solutions containing 0.1 M NaOH, nitramine formation increased when

294

the NO2 concentration decreased from 100,000 to 1,000 ppm.48 This excluded N2O4 as the major

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nitrating agent, because its formation should be favored at high NO2 concentrations (reaction 7).

296

Similarly, in MOR solutions (10 mM, pH 7.4), Cooney et al. observed that the positive

297

correlation between the NO2-MOR formation rate and NO2 concentration was less distinct at

298

NO2 concentrations above 25 ppmv.42 Additionally, when deuterium oxide was used as a solvent,

299

NO2-MOR formation was suppressed by 36%,42 consistent with the increased difficulty in

300

abstracting a deuterium to form the radical intermediate in the NO2-initiated radical pathway. C-

301

nitro compounds were detected when N-methylaniline was exposed to NO2, likely resulting from

302

electron delocalization of the aromatic amine radical prior to reaction with a second NO2.48

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HONO as a nitrosating agent: The role of nitrous acid (HONO) in nitrosation was

304

considered in atmospheric research. Hanst et al. and Pitts et al. monitored the transformation of

305

gaseous dimethylamine and diethylamine (0.5−3 ppmv) in the presence of 0.08−2 ppmv NO,

306

0.16−2 ppmv NO2, and ~30% humidity, and proposed that HONO was the contributing

307

nitrosating agent, based on the observation that nitrosamine formation was significantly slower

308

under conditions hindering HONO formation (reaction 8) such as low humidity or pre-dispersing

309

NOx in the chamber.49, 50 However, considering the low pKa (~3.4) of HONO and the basic pH of

310

amine solvents, HONO is unlikely to contribute significantly to nitrosamine formation in the

311

absorber.


Page 15 of 51


312

Lastly, it should be noted that some of the active nitrosating and nitrating agents in

313

biological systems are not expected to be relevant for CO2 capture systems and hence not

314

discussed here. These include nitrosyl halide, nitrosyl thiocyanate, and nitrous acidium ion, all of

315

which are only active under acidic conditions,45,

316

which requires NO oxidation by superoxide (O2•−).70-72

317

2.3

65-67

and peroxynitrite,68,

69

the formation of

Inhibitory Effects of CO2 and SO2

318

In a packed-bed absorber column reactor, NMOR and NO2-MOR accumulation rates

319

were 1.5-fold higher in the absence of CO2, although the rates remained approximately constant

320

over 1.6–12% CO2.33 Similarly, when purged with 100 ppmv NOx, unloaded MOR solution (5

321

M) formed 1.5 times more NMOR than MOR solution loaded with CO2 (0.36 C/N).35 CO2 may

322

inhibit nitrosation by either scavenging the nitrosating agent N2O346, 73 or forming a carbamate to

323

protect the amines from nitrosating agents.73,

324

scavenging by bicarbonate is 1.5×106 M-1 s-1 (reaction 6).46 Kirsch et al., however, used

325

NMR to argue that the inhibitory effect should be attributed to carbamate formation.74 Based on

326

DFT calculations, the reaction between N2O3 and dimethylcarbamic acid (i.e., protonated

327

carbamate) has a much higher energy barrier (117.65 kJ/mol) than that between N2O3 and

328

dimethylamine (6.32 kJ/mol).73 For primary and secondary amine solvents, the concentration of

329

amine carbamate far exceeds that of bicarbonate (e.g., the ratio of morpholine carbamate to

330

bicarbonate was 11.6 under 0.258 C/N loading at 40 °C).75 Therefore the carbamate inhibition

331

mechanism is likely to dominate. For tertiary amines, however, the bicarbonate inhibition

332

mechanism should be more important, because the absorbed CO2 primarily exists in the form of

333

bicarbonate and no carbamate is formed.76, 77

74

The second-order rate constant of N2O3


13

C


Page 16 of 51

334

SO2 also inhibits nitrosamine formation in the absorber. Chandan et al. observed 53%

335

lower NMOR formation from MOR in the presence of 84 ppmv SO2, and attributed it to the

336

reduction of NOx species by SO2.36 Alternative mechanisms, including the reaction between SO2

337

and amine, in a similar fashion as that between CO2 and amine,78 and the protonation of amines

338

upon absorption of the acid gas SO2, have not been considered.

339 340

3

DESORBER: Formation and Decomposition of Nitrosamines

341

The discussion in this section will focus on nitrosamines. Regarding nitramines, there

342

have been few mechanistic studies on their formation or decomposition in the desorber, but their

343

decomposition is anticipated to be facile given their low thermal stability. Indeed, increasing the

344

desorber temperature from 120 ºC to 140 ºC reduced the nitramine concentration in MEA solvent

345

by more than 75% in a pilot unit.21 Nitrate, a NOx hydrolysis product (reaction 10), was shown

346

not to contribute to nitramine formation in the solution phase.47

347

3.1

Nitrosamine Formation from Aqueous Nitrite: Kinetics and Mechanisms

348

In the desorber, which features elevated temperature and pressure, nitrosamines are

349

formed by reactions between amines and nitrite. Kinetic studies demonstrated that nitrosamine

350

formation rates are dependent on nitrite concentration, pH, CO2 loading, and, in some cases,

351

amine concentration.18, 20, 35, 79 The nitrite dependence of nitrosamine formation in the desorber is

352

different from that in acidic solutions reported by biological studies.45,

353

amines piperazine, MOR, diethanolamine, N-methylethanolamine, and N-(2-hydroxyethyl)

354

glycine, and the tertiary amine triethanolamine, the formation rates of nitrosamines at 100−145

355

ºC exhibited a first-order dependence on nitrite concentration.18, 20, 35, 79 The apparent first-order


60, 80-83

For secondary

Page 17 of 51


356

nitrosamine formation rate constants for piperazine, diethanolamine, and triethanolamine were

357

6.3×10-6 s-1, 1.0×10-6 s-1, and 6.9×10-7 s-1, respectively (100–120 ºC, 2.5–5 M amine, 0.2 C/N).18,

358

20

359

and secondary amines show a second-order dependence on nitrite concentration, attributed to

360

N2O3 formation being the rate-limiting step (reaction 9).45 Although nitrosation of tertiary

361

amines by nitrite at acidic pH exhibited a first-order dependence, it was attributed to the

362

nitrosonium ion (NO+) (reaction 11),84, 85 an agent unlikely to be important in the basic amine

363

solvents. It is worth noting that the nitrite concentrations used in current desorber studies range

364

from 0.6 to 100 mM,18,

365

level).86, 87 Research is needed to investigate how the nitrite concentration range affects reaction

366

mechanisms.

In contrast, at acidic pH and physiological temperature (37 ºC), nitrosation rates for primary

20, 79

slightly higher than those observed in field samples (0.2−9 mM

367

Regarding the dependence on pH, CO2 loading and amine concentration, a first-order

368

dependence on proton concentration and the CO2 loading of piperazine was observed for the rate

369

of 1-nitrosopiperazine formation from nitrite, but the dependence on piperazine concentration

370

was not reported.18,

371

concentration (1−5 M) in unloaded, self-buffered solutions.35 For triethanolamine, the total

372

nitrosamine formation was first-order with CO2 loading, and was not dependent on amine

373

concentration (0.5−2.5 M), but the pH dependence was not evaluated.20 The activation energy for

374

nitrosamine formation from piperazine, MOR, diethanolamine, and triethanolamine in the

375

presence of CO2 was 84 kJ/mol, 101 kJ/mol, 42 kJ/mol, and 74 kJ/mol, respectively.18, 35, 79 The

376

kinetic models suggest that carbamic acids of secondary amines are the precursors of

377

nitrosamines under desorber conditions.18,

378

aqueous phase, the rate-limiting step for the formation of N-nitrosodimethylamine from

79

The NMOR formation rate was 0.25-order dependent on MOR

79

Similarly, DFT calculations indicated that, in the



Page 18 of 51

379

dimethylamine in the presence of CO2 was the formation of a four-membered ring transition state

380

from carbamic acid and nitrite (Scheme 1e),73,

381

alternative mechanism was proposed involving an initial reaction between CO2 and NO2⎻ to form

382

ONOCO2⎻, which subsequently serves as the nitrosating agent for amines (Scheme 1f),73 but it

383

does not account for the pH dependence of nitrosamine formation. Although NO+ was also

384

considered a nitrosating agent in the desorber,35 the absence of CO2 in the experiments and the

385

lack of evaluation on pH dependence renders it inconclusive.

88

with an energy barrier of 188 kJ/mol.73 An

386

In earlier biological studies, formaldehyde and other carbonyl compounds were shown to

387

catalyze nitrosation of secondary amines by nitrite at neutral and alkaline pH.45, 80, 89, 90 As shown

388

in Scheme 1g, a carbinolamine is formed after the nucleophilic attack of a secondary amine on

389

formaldehyde, the rate-limiting step with an energy barrier of 113−167 kJ/mol.91 The protonated

390

carbinolamine then dehydrates to form an iminium ion,90,

391

nitrite to form a nitrosamine.89-91 At pH above 7.5 and room temperature, the addition of 50 mM

392

formaldehyde increased the yield of nitrosamines from diethylamine from non-detectable to

393

approximately 0.25% (reaction time 17 h).89 The formaldehyde-catalyzed formation of NMOR

394

and N-nitrosodiethylamine from MOR and diethylamine, respectively, both exhibited a first-

395

order dependence on nitrite.90 Acetaldehyde is expected to exhibit similar but lower catalytic

396

activity.91 Aldehydes have been reported as oxidative degradation products of MEA92-94 and have

397

been detected at hundreds of ppbv levels in the exhaust gas of pilot- and industrial-scale CO2

398

capture plants,95, 96 but the role of aldehydes for nitrosamine formation under desorber conditions

399

has not been evaluated.

400

3.2

91

Nitrosamine Decomposition


which subsequently reacts with

Page 19 of 51


The high temperatures used in desorbers can also promote nitrosamine decomposition.19,

401 402

21

403

140 ºC resulted in a 50% reduction in total nitrosamines.21 The decomposition of four

404

nitrosamines, NMOR, 1-nitrosopiperazine, NDELA, and N-nitroso-(2-hydroxyethyl) glycine,

405

followed first-order kinetics,19,

406

respectively;19, 35 correspondingly, 1-nitrosopiperazine decomposed the fastest among these four

407

nitrosamines, with a half-life of 6.7 h (150 ºC, 4.9 M piperazine, 0.1 C/N).19

In a pilot unit using a 30% MEA as solvent, increasing desorber temperature from 120 ºC to

35, 97

with activation energy of 131, 94, 106, and 112 kJ/mol,

408

Higher solution pH and the presence of strong base (e.g., solvent amines) promoted

409

1-nitrosopiperazine decomposition, which was attributed to a base-catalyzed pathway.19

410

Base-catalysis could explain the faster nitrosamine decomposition in diethanolamine (pKa 8.90)

411

solvent than in triethanolamine (pKa 7.74) solvent, with half-lives of 91 and 231 h, respectively

412

(120 ºC, 2.5 M amine, 0.2 C/N).20 In the absence of high concentrations of amines, thermal

413

decomposition of nitrosamines also followed first-order kinetics across the pH range 2.2−12.5 at

414

110 ºC.98

415

Nitrosamine decomposition can differ regarding pH dependence and product distribution,

416

but denitrosation accounts for the majority of nitrosamine decomposition. After heating at 150 ºC

417

for two days, all nitroso-equivalents in 1-nitrosopiperazine were removed from the aqueous

418

phase as gaseous products; the major decomposition product 2-piperazinol accounted for half of

419

the 1-nitrosopiperazine degraded.19 At a lower temperature (110 ºC) and without amines present,

420

nitrite evolved over the course of 5 days at pH 8.5−12.5 as N-nitrosopyrrolidine degraded, but it

421

accounted for only 10–50% of the degraded N-nitrosopyrrolidine.98 Decarboxylation of N-

422

nitrosoproline and N-nitrososarcosine was insignificant.98 Nitrosamines featuring a carboxyl

423

group adjacent to the amino nitrogen (e.g., N-nitrososarcosine and N-nitrosoproline) decomposed



Page 20 of 51

424

fastest at acidic pH (2.2−4), while cyclic nitrosamines (e.g., N-nitrosopyrrolidine) decomposed

425

fastest at alkaline pH.98 Nitrosamines featuring β-hydroxy groups can undergo base-induced

426

fragmentation to form an aldehyde or ketone, a vinylnitrosamine, and smaller alkylnitrosamines;

427

for example, NDELA decomposition products include formaldehyde, methylvinylnitrosamine,

428

NMOR, and N-nitrosodimethylamine.99 The fragmentation rate increases with base concentration

429

and steric hindrance of the β-carbon.99 A group of unstable nitrosamines, cyclic α-hydroxy-N-

430

nitrosamines (e.g, 1-nitroso-pyrrolidin-2-ol), decay to form diazonium ions,100 but these

431

nitrosamines are likely too unstable to be present at any significant concentration in the solvent.

432 433

4

Influence of Solvent Amine Structure on Nitrosamine and Nitramine Formation

434

4.1

Effects of Amine Order

435

The effects of amine structure on NOx-driven nitrosamine formation in the absorber were

436

systematically investigated by Dai and Mitch in a lab-scale packed-bed column reactor.16 By

437

employing a total nitrosamine analysis method, which indiscriminately measures all N-nitroso

438

functional groups, the authors showed that the order of the amino group was a critical

439

characteristic determining the nitrosamine formation potential of amines across alkanolamines,

440

straight-chain and cyclic diamines, and amino acids.16 Secondary amines diethanolamine,

441

dimethylethylenediamine, piperazine, and sarcosine formed nitrosamines the fastest, closely

442

followed (within a factor of three) by tertiary amines triethanolamine, N-methyldiethanolamine,

443

tetramethylethylenediamine, and 1,4-dimethylpiperazine, while primary amines MEA,

444

ethylenediamine, and glycine showed approximately two orders of magnitude lower reactivity.16

445

Within primary amines, structural characteristics including steric hindrance of the amino group


Page 21 of 51


446

in alkanolamines, alkyl chain length of straight-chain diamines and amino acids, and β-

447

substituents of the amino group did not significantly affect nitrosamine formation.16 Jackson and

448

Attalla also reported higher reactivity of secondary amines than primary and tertiary amines:

449

after purging amine solutions with 8000 ppmv NO and 5% O2 for 15 hours, strong nitrosamine

450

signals were observed by mass spectrometry for the secondary amines piperazine,

451

diethanolamine, and 2-piperidinemethanol, but not for the primary amines MEA and

452

2-amino-2-methyl-1-propanol or the tertiary amine N-methyldiethanolamine.101 Nitrosamine

453

formation kinetics has not been thoroughly investigated in amine mixtures, but some preliminary

454

findings suggest complex interactions may occur. For example, total nitrosamines formed in a

455

mixture of 2.5 M tertiary amine N-methyldiethanolamine and 0.5 M secondary amine piperazine

456

was only a third of that formed in a 0.5 M piperazine solution.16

457

In the desorber, the order of the amino group was also a determining factor for

458

nitrosamine formation.16 Secondary amines exhibited the highest nitrosamine formation rate,

459

followed by tertiary amines (within a factor of eight).16, 20 Nitrosamine formation from primary

460

amines was more than three orders of magnitude lower than from secondary amines.16 The yields

461

of nitrosamines based on nitrite loss vary among amines. The conversion of nitrite to 1-

462

nitrosopiperazine followed 1:1 stoichiometry,18,

463

10−60% of nitrite consumption for secondary and tertiary alkanolamines (e.g., diethanolamine,

464

2-(methylamino)ethanol, triethanolamine, and 2-(diethylamino)ethanol).20

79

but the yields of nitrosamines were only

The effects of amine structure on nitramine formation have not been investigated in the

465 466

absorber or desorber.

467

4.2

468

Tertiary Amines

Mechanisms of Nitrosamine and Nitramine Formation from Primary, Secondary, and



Page 22 of 51

469

Nitrosamine formation can be considered as a two-step process: an initial interaction

470

between an amine and a nitrosating agent to form a reactive intermediate, followed by further

471

transformation of the reactive intermediate to form a nitrosamine. Because amine order primarily

472

affects the second step of the process, the dependence of nitrosamine formation on amine order

473

in CO2 capture systems is similar to those observed in earlier biological studies, and can be

474

explained by some common mechanisms. For secondary amines, stable nitrosamines form in one

475

step upon amines reacting with nitrosating agents.45 In contrast, primary amines, upon interaction

476

with nitrosating agents, form unstable primary nitrosamines, which decay to a wide variety of

477

deamination products via a carbonium ion intermediate (Scheme 1h).65, 102 Because secondary

478

amines can form from the carbonium ion, nitrosamines can form via the subsequent nitrosation

479

of these secondary amines.102 The yields of nitrosamines from primary amines in acidic nitrite

480

solutions (i.e., N-nitrosodimethylamine from methylamine, N-nitrosodibutylamine from n-

481

butylamine, N-nitrosopiperidine from 1,5-pentanediamine, and N-nitrosopyrrolidine from 1,4-

482

diaminobutane) were between 0.002% and 3% based on amine consumption.103, 104 This pathway

483

is applicable for primary amines in both absorber and desorber.

484

For tertiary amines, three different nitrosation pathways were proposed for acidic nitrite

485

solutions, but only two apply to the desorber of CO2 capture systems. Tertiary amines undergo

486

nitrosative dealkylation prior to forming a nitrosamine.84, 85, 105, 106 As shown in Scheme 1i, the

487

three nitrosative dealkylation pathways proposed by Loeppky and colleagues all involve two

488

common intermediates: a nitrosammonium ion (A) after the initial nitrosation, and an iminium

489

ion (B) prior to dealkylation.84, 85, 105 In pathway I, the primary pathway for tertiary alkylamines

490

in acidic nitrite solutions, A undergoes heterolytic elimination of nitroxyl (NOH) to form B,

491

accompanied by N2O evolution.84,

105

Pathways II and III, proposed for N,N-dialkyl aromatic


Page 23 of 51


492

amines, both involve homolysis of A to form an amine radical cation (R2ArN•+) intermediate

493

C.84, 85 C either loses a proton from the N-bound alkyl group to a Lewis base, followed by rapid

494

oxidation to B (pathway II) or undergoes NO2-mediated α-H atom abstraction to form B

495

(pathway III).84, 85 As the tertiary amines used in CO2 capture lack aromatic functionality, the

496

NOH elimination pathway (I) would be expected to dominate in CO2 capture systems; however,

497

this was not supported by experimental results. Under simulated desorber conditions, Yu et al.

498

observed a regioselectivity of dealkylation contrary to that expected from pathway I.20 For

499

tertiary amines 2-(dimethylamino)ethanol and N-methyldiethanolamine, cleavage of the 2-

500

hydroxyethyl group was twice more likely than demethylation,20 but pathway I is associated with

501

preferential cleavage of the smaller substituted group.84, 105 Overall, under desorber conditions,

502

pathway II, and possibly III, are likely preferred over pathway I for nitrosamine formation from

503

tertiary amines.20 This proposition is also consistent with DFT calculations for nitrosative

504

cleavage of trimethylamine in both gas and aqueous phases, which suggested that the homolysis

505

of nitrosammonium ion (pathways II and III) dominated over NOH elimination (pathway I).106

506

Tertiary amine nitrosation mechanisms have not been investigated under absorber conditions.

507

The influence of amine order on nitramine formation is different from that on nitrosamine

508

formation. Unlike primary nitrosamines, primary nitramines are sufficiently stable to be detected

509

in aqueous solutions.13, 57, 107 For example, the hydrolysis lifetimes of 2-(nitroamino)ethanol and

510

dimethylnitramine were more than three months across the pH range 5–9.13 Because the low net

511

nitrosamine formation from primary amines is attributed to the instability of primary

512

nitrosamines, it is reasonable to expect that the stability of primary nitramines would render

513

nitramine formation from primary and secondary amines to be similar. Secondary amines form

514

nitramines directly.47, 48 Tertiary amines, due to their higher steric hindrance, and requirement to



Page 24 of 51

515

undergo dealkylation, are likely to have lower nitramine formation potential than primary and

516

secondary amines. These trends, however, need to be verified in CO2 capture systems.

517

Additionally, a total nitramine method capable of capturing all compounds featuring the N-nitro

518

group is needed for the comparison of nitramine formation from different amines.

519

4.3

Other Structural Characteristics

520

Substituents and steric hindrance of the amino group were shown to affect nitrosamine

521

and nitramine formation by computational, biological and atmospheric chemistry studies,57, 61, 73

522

but their effects have not been observed under CO2 capture scenarios. Using DFT calculations,

523

Sun et al. and Zhao et al. showed that electron-donating substituents would decrease the energy

524

barrier of N2O3-initiated and NO2-initiated nitrosation reactions in the gas phase, respectively.61,

525

62

526

primary alkanolamines in the absorber,16 were reported to slow reactions potentially relevant to

527

the desorber, including aldehyde-catalyzed nitrosation (Scheme 1g)89,

528

formation from tertiary alkanolamines (Scheme 1i),20 both involving a “crowded” transition

529

state.

530

4.4

Higher steric hindrance, while having little impacts on the formation of nitrosamines from

90

and nitrosamine

Distribution of Specific Nitrosamines: Importance of Amine Degradation

531

Mass spectrometry-based analyses have been employed in parallel with total nitrosamine

532

analysis to identify the distribution of specific nitrosamine products,21, 22, 37, 108 which reflects the

533

nitrosation mechanisms of amines. Because secondary amines form nitrosamines directly upon

534

nitrosation in both the absorber and desorber, they feature the least diverse nitrosamine products.

535

When a MOR solution was purged with 25 ppmv NO and NO2, NMOR accounted for all of the

536

total nitrosamine signal.25 Under desorber conditions, 1-nitrosopiperazine accounted for more


Page 25 of 51


537

than 99% of nitrite loss in piperazine solutions.18 In comparison, nitrosation of primary and

538

tertiary amines forms a variety of specific nitrosamines, as a result of the respective deamination

539

and dealkylation steps prior to nitrosamine formation (Schemes 1h and 1i).20,

540

alkanolamines, nitrosative dealkylation under desorber conditions is more likely to remove 2-

541

hydroxylethyl or ethyl groups than a methyl group.20 For example, nitrosamines formed from

542

dimethylethanolamine were half N-nitrosodimethylamine and half N-nitrosomethylethanolamine;

543

those formed from N-methyldiethanolamine were 66% N-nitrosomethylethanolamine and 17%

544

NDELA.20 It is worth noting, however, that minor nitrosamine products, if volatile, may

545

dominate emissions. In a packed-bed column unit operated with sodium sarcosinate solvent, N-

546

nitrosodimethylamine formation was at least ten times slower than N-nitrososarcosine in the

547

absorber, but N-nitrosodimethylamine accounted for 30% of the total nitrosamines accumulated

548

in the washwater whereas N-nitrososarcosine was not detected, indicating significant emission of

549

N-nitrosodimethylamine from the absorber.16 N-Nitrosodimethylamine was also detected as a

550

minor specific nitrosamine from primary amine MEA and several tertiary alkanolamines. 20, 21

551

Impurities in commercial solvents, such as diethanolamine in MEA solvents,17 can also

552

contribute to nitrosamine formation.

21

For tertiary

553

Because amines are the active agent for CO2 absorption, their nitrosation is also

554

considered as “nitrosative degradation” of amines. In addition to nitrosative degradation, amines

555

undergo oxidative degradation in the absorber and thermal degradation in the desorber, forming a

556

spectrum of nitrosamine precursors and further complicating the distribution of specific

557

nitrosamines. Comprehensive reviews of amine degradation are available elsewhere;109, 110 here

558

we focus on the findings relevant to the formation of nitrosamines. Upon oxidative and thermal

559

degradation, MEA and 2-amino-2-methyl-1-propanol, the two most studied primary amines,



560

form two groups of products with distinct nitrosamine formation potential. The first group

561

features a secondary or tertiary amino group in polyamines or derivatives of piperazinones,22, 111

562

and is anticipated to contribute to the majority of nitrosamine formation. The second group,

563

including derivatives of amides, oxazolidinones, imidazolidinones, and ureas, features a carbonyl

564

group adjacent to the nitrogen atom.111-113 This latter group is expected to be inactive towards

565

nitrosation in the absorber due to the strong electron withdrawing property of the carbonyl group,

566

but may form small amounts of nitrosamines in the desorber via pathways similar to those of

567

carbamates (Scheme 1e). Among tertiary amines, N-methyldiethanolamine is the most studied

568

for oxidative and thermal degradation. Many of the N-methyldiethanolamine degradation

569

products retain tertiary amine functionality, but secondary amines such as diethanolamine and

570

N-methylethanolamine are also formed.112,

571

nitrosamine formation potential than the parent N-methyldiethanolamine.20

114

Page 26 of 51

These secondary amines would exhibit higher

572

Due to the concurrence of multiple amine degradation pathways, the distribution of

573

nitrosamine products cannot be predicted by nitrosative, thermal, or oxidative degradation alone.

574

For example, in a pilot unit operated with 30% MEA, only 2% of total nitrosamines were

575

NDELA,21 the primary specific nitrosamine anticipated from nitrosative degradation of MEA

576

(Scheme 1h). Under simulated absorber conditions, dissolved metals iron and copper both

577

promoted MEA oxidation,115 but only copper promoted the formation of nitrosamines.116

578

Moreover, many of the current thermal/oxidative degradation studies employ extreme

579

conditions,78, 117-119 and therefore may not capture the degradation intermediates that are reactive

580

nitrosamine precursors. For example, N-(2-hydroxyethyl)glycine was not reported as a major

581

product from MEA degradation,22,

582

56% of total nitrosamines in a pilot unit.21 These findings highlight the importance of field

108-110

but N-nitroso-(2-hydroxyethyl)glycine accounted for


Page 27 of 51


583

sampling of pilot- and industrial-scale systems. In addition, degradation experiments in the

584

presence of nitrite would be valuable for exploring the interplay between nitrosative and

585

thermal/oxidative degradations. In the presence of nitrite (5000 ppm), thermal degradation of

586

MEA was enhanced by five-fold, and diethanolamine was the third most abundant degradation

587

product, which was not observed in the absence of nitrite.120

588

The distribution of specific nitramines in CO2 capture systems has not been investigated,

589

attributable to the absence of a total nitramine analysis method, which prohibits the

590

quantification of the contribution of specific nitramines to the total pool of nitramines.

591

Additional challenges include the difficulty in obtaining analytical standards for specific

592

nitramines and the low instrument sensitivities to nitramines due to their low thermal stability.

593 594

5

WASHWATER UNIT: Mass Transfer and Formation in situ

595

Washwater units are used to recover amines and control emissions of volatile

596

contaminants, including nitrosamines and nitramines.25, 121-124 Accumulation of nitrosamines in

597

the washwater has been documented for pilot-scale and laboratory-scale systems. In a pilot

598

system operated with 35% MEA, 0.042 µM NDELA and 0.73 µM total nitrosamines were

599

detected in the washwater; when the system was operated with a mixture of 25% 2-amino-2-

600

methyl-1-propanol and 15% piperazine, 59 µM total nitrosamines were detected, reflecting the

601

higher nitrosamine formation from the secondary amine piperazine than from the primary amine

602

MEA.25 In a laboratory-scale reactor, nitrosamine accumulation in the washwater unit was

603

approximately ten times slower than in the absorber solvent for many secondary and tertiary



604

amines.17,

25, 121, 125

605

nitrosamine in the washwater.121

Page 28 of 51

In the same reactor, nitramine accumulation was 200 times slower than

606

Two pathways for nitrosamine accumulation in the washwater have been identified.121, 125

607

First, the nitrosamines accumulating in the solvent, via either absorber or desorber reactions, can

608

transfer to the washwater with the absorber exhaust gas. Second, amines accumulating in the

609

washwater can react with residual NOx in the absorber exhaust to form nitrosamines in situ (i.e.,

610

within the washwater unit).121,

611

nitrosopiperazine in the washwater when 100 mM of piperazine was spiked in the washwater of

612

a system operated with a MOR solvent.125 When MOR or diethanolamine were used as solvent

613

amines for flue gases containing equimolar 25 ppmv NO and NO2 or 15 ppmv NO and 3 ppmv

614

NO2, in situ formation accounted for more than 80% of nitrosamine accumulation in the

615

washwater after 6 h.121,

616

quadratic time dependence of the nitrosamine concentration in the washwater (i.e., nitrosamine

617

accumulation accelerated over time due to accumulation of amines in the washwater).121 High

618

flue gas NOx concentration promotes nitrosamine accumulation through both the mass transfer

619

and in situ formation pathways. The accumulation rate of NMOR in the washwater associated

620

with MOR solvent exhibited a first-order dependence on NO and NO2 gaseous concentrations,

621

respectively.33 Although there are no detailed reports on nitramine accumulation in the

622

washwater, similar pathways (transfer from the absorber and in situ formation) are anticipated.

125

125

The latter pathway was confirmed by the detection of 1-

Modeling showed that the in situ formation pathway resulted in a

623

The volatility of solvent amines and their associated nitrosamines heavily influences

624

nitrosamine accumulation in the washwater.16 For example, MOR and NMOR are 4−6 orders of

625

magnitude more volatile than diethanolamine and NDELA, respectively, as a result of the

626

alcohol functional groups in the latter.126 Accordingly, nitrosamine accumulation in the


Page 29 of 51


627

washwater associated with MOR resulting from transfer and in situ formation was 3200 and

628

1400 times higher than that with diethanolamine, respectively.121 The accumulation of

629

nitrosamines in the washwater associated with triethanolamine and MEA was comparable,

630

despite the ten times higher reactivity of triethanolamine.16 This can be explained by the higher

631

volatility of MEA, which contains only one alcohol group, promoting its transfer from the

632

absorber to the washwater and allowing subsequent reactions with NOx in the washwater to form

633

nitrosamines.126 The transfer mechanism for nitrosamines from the solvent or washwater to the

634

gas phase is unclear. The aerosol phase is known to dominate amine emissions,127-129 but a field

635

study indicated that nitrosamines were only detected in the gas phase but not in aerosol

636

droplets.130 Further study on the abundance of nitrosamines and nitramines in the aerosol phase

637

is needed.

638 639

6

Mitigation Strategies for Nitrosamines and Nitramines in CO2 Capture Systems

640

Nitrosamines and nitramines from CO2 capture systems enter the environment via stack

641

emissions and disposal of spent solvents. Spent solvents are typically incinerated or undergo

642

biological treatment.131-133 Compared to the disposal of spent solvents, stack emissions

643

continuously introduce contaminants to the environment and impact downwind communities.

644

Strategies for controlling stack emissions of nitrosamines and nitramines include: 1) destroying

645

them in the desorber, 2) enhancing washwater performance, and 3) preventing their formation

646

through NOx removal prior to CO2 capture and/or optimization of amine solvents.

647

High temperature desorption was shown to facilitate nitrosamine decomposition,20, 79 but

648

it is applicable only to amines with high thermal stability such as piperazine, and it comes with



Page 30 of 51

649

the penalty of increasing energy cost. Metal-based heterogeneous catalysts were also developed

650

to destroy nitrosamines under desorber conditions, but their use may be limited due to the need

651

for hydrogen gas.134 Additionally, catalyst deactivation is anticipated given the extremely

652

complex solvent matrix.135

653

An in-line washwater treatment system was proposed to enhance the physical absorption

654

efficiency of washwater units by continuously removing the nitrosamines and nitramines

655

accumulated in the washwater. Because nitrosamines, and presumably nitramines, accumulate in

656

the washwater through two pathways, transfer from the absorber and in situ formation, the

657

treatment system employed ultraviolet (UV) and ozone to address each pathway respectively.16,

658

121

659

mJ/cm2 UV incident fluence from a medium pressure mercury lamp, the accumulation of NMOR

660

and NO2-MOR was reduced by an order of magnitude over the course of 6 h in the washwater

661

associated with MOR solvent.121 More importantly, the application of UV eliminated the

662

acceleration in nitrosamine accumulation in washwater, suggesting that its benefits could

663

increase over time.119 Ozonation, although it destroyed amines to suppress in situ formation, was

664

ineffective in further reducing the accumulation of nitrosamines and nitramines, because

665

ozonation products absorbed UV and reduced the photolysis efficiency;121 however, ozonation

666

does have the potential to control amine emissions in the final exhaust and thereby minimize

667

formation of nitrosamines and nitramines by reactions between amines and ambient NOx. The

668

efficacy of in-line washwater treatment has not been evaluated in parallel with gas phase

669

monitoring.

Nitrosamines and nitramines undergo direct photolysis under UV.136,

137

With 272−537

670

Removing NOx prior to CO2 capture and developing solvents of low nitrosamine and

671

nitramine formation potential offer two promising pollution prevention strategies. Because


Page 31 of 51


672

nitrosamine formation in the absorber increases with NO (0–22.5 ppmv) and NO2 (0–15 ppmv),

673

and nitramine formation increases with NO2 (0–15 ppmv),33, 34 advanced NOx removal to sub

674

ppmv level is not expected to exhibit diminished returns. NOx removal also reduces nitrosamine

675

formation in the desorber by forming less nitrite in the solvent,20, 33, 79 and minimizes the in situ

676

formation of nitrosamines in the washwater.121 It should be recognized, however, that advanced

677

NOx removal may add to the cost of CO2 capture.

678

Developing solvents of low nitrosamine and nitramine formation potential has not been

679

considered systematically in conjunction with efforts to improve the CO2 capture performance of

680

solvents.

681

N-methyldiethanolamine were proposed to take advantage of the fast absorption rate of MEA

682

and the low regeneration energy and high capacity of N-methyldiethanolamine.6, 138 This mixture

683

may have the added benefit of lowering nitrosamine formation due to the net consumption of

684

nitrosating agents through the rapid decay of the primary nitrosamine, but it has not been

685

evaluated. Amino acid salts (e.g., potassium sarcosinate) have attracted attention due to their low

686

volatility compared to amines,139-141 but nitrosamine and nitramine formation from amino acid

687

salts has rarely been considered. Additionally, as discussed previously, the extremely volatile

688

nitrosamine N-nitrosodimethylamine was detected in the washwater associated with sodium

689

sarcosinate,16 suggesting that a detailed product analysis is needed even for seemingly non-

690

volatile amino acids. Secondary amines, due to their relatively high nitrosamine formation

691

potential, should be avoided unless other control strategies are employed.

For

example,

mixtures

of

primary

amine

692 693

7

Future Research Needs


MEA

and

tertiary

amine


Page 32 of 51

694

This review focuses on the current knowledge on the formation and decomposition of

695

nitrosamines and nitramines in the absorber, desorber, and washwater units of CO2 capture

696

systems. Despite advances in fundamental understanding, a quantitative model is still lacking

697

that is able to predict the concentrations of total and specific nitrosamines and nitramines in the

698

solvent and in the exhaust gas. Such a model would use inputs including flue gas composition,

699

solvent composition (e.g., type and concentration of amines), and other operating conditions

700

(e.g., CO2 loading and temperature), and take into consideration the effects of amine degradation

701

and the interaction between the absorber, desorber, and washwater units.

702

Three research gaps are hindering the development of such a predictive model. First,

703

uncertainties remain in the nitrosation and nitration mechanisms for CO2 capture systems.

704

Specifically, there is no definitive conclusion regarding whether NO2 or N2O3 initiates

705

nitrosation in the absorber. The two pathways not only give rise to different kinetics, they also

706

affect the relative rates of nitrosation to other NOx reactions, including but not limited to

707

hydrolysis, dimerization, dissociation, and oxidation (reactions 3–4, 7–10) in both the gas and

708

solvent phases.18, 33, 35, 142, 143 In addition, because NOx species originate from the flue gas while

709

amines are present in the solvent phase, mass transfer of NOx into the solvent phase also affects

710

system behavior. Many laboratory-scale CO2 capture reactors were not able to quantify NOx

711

reactions or mass transfer, which may confound nitrosation kinetics, limiting their use for

712

probing reaction mechanisms. The use of a well-characterized wetted wall column reactor by

713

Fine et al. was to date the best attempt to quantify NOx mass transfer in parallel to nitrosamine

714

formation,34, 97 but full consideration of NOx reactions and quantification of nitrosamines are still

715

needed.


Page 33 of 51


716

Second, the distribution of specific nitrosamines and nitramines is largely unknown or

717

determined on a case-by-case basis. Such information is critical for determining reaction

718

mechanisms, estimating emissions, and assessing the environmental fate and impacts of these

719

compounds. For nitramines, a total nitramine method is lacking, and targeted analysis has been

720

limited to a few nitramines with available analytical standards.21 Even for nitrosamines, a high

721

percentage of total nitrosamines often remain unidentified.21, 25 Non-targeted analysis based on

722

high-resolution mass spectrometry is a promising tool,22, 144 but method development is needed

723

for the challenging solvent matrix. Washwater may be a good starting point for non-targeted

724

analysis due to its simpler matrix and preferential enrichment of volatile nitrosamines and

725

nitramines. Trace nitrosamines and nitramines, if sufficiently volatile, can dominate emissions,

726

but are often overlooked in current studies due to analytical difficulties. Improving the sensitivity

727

and frequency of gas phase sampling will contribute to determining important nitrosamines and

728

nitramines with high emission potential.

729

Third, there is a disconnect among studies on the individual units of the CO2 capture

730

systems. As amine solvents circulate between the absorber and desorber, the NOx absorption

731

efficiency of the absorber and oxidative degradation of amines affect the amount of nitrosating

732

agent (i.e., nitrite) and nitrosamine precursors in the desorber; in turn, thermal degradation of

733

amines in the desorber alters the profile of nitrosamine and nitramine precursors in the absorber.

734

Studies with recirculating absorber and desorber units are needed to capture these effects.

735

Washwater units allow additional reactions between accumulated amines and residual NOx in the

736

flue gas, a pathway most relevant for amines and their degradation products of intermediate

737

volatility–those sufficiently volatile to be emitted from the absorber yet not so volatile as to

738

escape the washwater. The emission pathways for nitrosamines and nitramines remain unclear.



Page 34 of 51

739

It is worth noting that in addition to formation and decomposition within CO2 capture

740

systems, other critical inputs are needed for a comprehensive risk assessment regarding

741

nitrosamines and nitramines, including their environmental fate in the atmosphere and surface

742

water13, 145 and their toxicity.146, 147 Their concentration in the CO2 product stream also should be

743

considered before the CO2 is considered for certain commercial uses, such as the food industry.

744

Additionally, the emission of amines also needs to be considered, because they may react with

745

ambient NOx to form nitrosamines and nitramines in the atmosphere.12, 13, 27 An interdisciplinary

746

effort is needed to address the challenge of nitrosamines and nitramines in CO2 capture systems.


Page 35 of 51

747


Figure 1. Amine-based CO2 capture system.

748 749 750 751



752

Page 36 of 51

List of Acronyms Acronym

Definition

MEA

Monoethanolamine

MOR

Morpholine

NMOR

N-Nitrosomorpholine

NDELA

N-Nitrosodiethanolamine

NO2-MOR

N-Nitromorpholine

753


Page 37 of 51

754


Table 1. NOx-related reactions in CO2 capture systems Reaction

Rate Constants (aqueous)a

Reference

kF = 1.1×109 (M −2 s−1 )

148

(1) (2) (3)

kB = 8.03×104 (s −1 )

(4)

ln kF = −876.1/ T + 17.52 (M −2 s −1 )

149

(5)

ln kB = −1912.4 / T + 13.79 (s−1 )

150

kCl = 1.44 ×105 (M −1s−1 ) −

(6)

− 2

N2O3 + X → XNO + NO −

−

− 3

2− 4

− 4

( X = Cl , HCO , HPO , H2PO )

kHCO- = 1.49 ×106 (M −1s−1 )

44, 46

3

kphosphate = 6.4 ×105 (M −1s −1 )

(7)

755

kF = 4.5 ×108 (M −1s−1 ) 3

151

−1

kB = 6.9 ×10 (s )

(8)

kF = 8 ×109 (M −1s−1 )

152, 153

(9)

ln kF = −1912.4 / T + 13.79 (s−1 )

150

(10)

ln kF = −9526.7 / T + 37.62 (s −1 )

154

(11)

N/A

a

Constant reported at 25 °C if temperature relationship is not shown.

756 757



758 759

Table 2. Solvent and washwater characteristics in pilot- and full-scale MEA-based CO2 capture systems Description Temperature

Absorber: 40–60 °C 10 Desorber: 100–140 °C 10

30 °C

10−13 51

~10 b

Amine Concentration

2.5–8 M 9

~10 mM

Nitrate Concentration Nitrosamine Concentration Aldehyde Concentration

762

Washwater a 25

Solvent

Solution pH Nitrite Concentration

760 761

Page 38 of 51

0.1–10 mM

22, 39

100−20 µM

20−200 mM

22, 39

~20 µM

50−3000 µM

21, 22

1−100 µM

0.6–2 mM

39

a

150−200 µM

Washwater samples collected from the Aminox pilot reactor at Statoil ASA without desorber unit; residence time of washwater 10 hr. b Acidic washwater (e.g., pH 3−7) has been proposed to maximize absorption of amines.155


Page 39 of 51

763 764 765 766 767 768


Scheme 1. Pathways for nitrosamine and nitramine formation. (a–g) Formation of Nnitrosodiethanolamine (NDELA) and N-nitrodiethanolamine from the secondary amine diethanolamine. (h) Nitrosamine formation from the primary amine monoethanolamine (MEA). (i) Pathways for nitrosamine formation from tertiary amines. (a) NO2-initiated radical pathway for nitrosamine and nitramine formation

769 770 771

(b) Formation of N2O3 isomers

772 773 774

(c) N2O3 pathway for nitrosamine formation

775 776 777

(d) N2O4 pathway for nitramine formation

778 779 780

(e) Nitrosamine formation from the reaction between carbamic acid and NO2−

781 782 783



784

(f) Nitrosamine formation from the reaction between amine and ONOCO2−

785 786 787 788

(g) Nitrosamine formation from the reaction between amine and NO2− catalyzed by formaldehyde

789 790 791

(h) Decay of primary nitrosamine

792 793 794

(i) Nitrosamine formation from tertiary amine

795 796


Page 40 of 51

Page 41 of 51


797

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48. 49.

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52.

53.

54. 55. 56.

57.

58.

59.

60.

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Nitrosamines and Nitramines in Amine-Based Carbon Dioxide

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