Corso di laurea magistrale in:
Chimica Fisica
Reversible CO
2capture by ethylene
glycol/potassium hydroxide-based liquid sorbents
Relatore: Prof. Gianluca Ciancaleoni
Controrelatore: Dr. Marco Taddei
Candidato: Marcello Costamagna
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Index
Introduction ... 3
Carbon Capture and Storage (CCS) ... 4
Methods of regeneration ... 10
Storage and reuse ... 10
Deep Eutectic Solvents and the Hole Model ... 13
The Hole Model ... 14
Process of discovery ... 18
Experimental section ... 22
Synthesis and carbonation ... 22
NMR ... 22 Electrochemical analysis ... 23 Thermogravimetric Analysis ... 23 Conductivity ... 23 Viscosity ... 24 Density ... 24 Results ... 25 Ethylene Glycol ... 25 KOH/EG ... 27 KOH/H3BO3/EG ... 30 KOH/EG + CO2 ... 33 KOH/H3BO3/EG + CO2 ... 36 Electrochemical results ... 43 Walden Plot ... 46 Melting Points ... 47 Discussion ... 48
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Conclusions ... 64 Bibliography ... 65
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Introduction
Climate change is one of the most important challenges humankind must face in the 21st century. CO
2 is the major greenhouse gas linked to climate change and its level in our atmosphere have been steadily increasing since the industrial revolution, jumping from an average of 250 ppm to more the 400 ppm, due to anthropic emission.
Figure 1. Partial history of global CO2 levels in the atmosphere1
We are now in a transition period from an economy based on the exploitation of fossil fuels to the plethora of available renewables and green alternatives. The goal set by the Intergovernmental Panel for Climate Change (IPCC) is the mitigation of climate change in order to keep the raising global temperature below 2°C within this century. Unfortunately this transition is not happening nearly as fast as it should, and in the IPCC’s 5th Assessment Report (AR5), published in the 2014,
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it has been highlighted the essential role that Carbon Capture and Storage (CCS) will have to fulfil in this century, alongside renewable energies.2 This was particularly made clear in an Integrated Assessment Model (IAM) in 2014,3 specifically targeted at the role of CCS in long-term climate change mitigation scenarios where 27 energy models for the century were compared. While the amount of CO2 captured varied widely between the models (from 600 to 3050 GtCO2 within 2100), they all consistently drew the conclusion of the impossibility of achieving the 2°C target without a proper development and implementation of CCS strategies. The reality is that the world today relies on fossil fuels for over 85% of its energy use. Changing that picture dramatically will take time. CCS thus offers a way to get large CO2 reductions from power plants and other industrial sources until cleaner, sustainable energy sources and technologies can be widely deployed as suggested by the new official European strategies.4 Other studies5,6 also indicate that by 2030 and beyond CCS will be a major component of a cost-effective portfolio of emission reduction strategies. Even though CCS is recognised as being vital to least cost pathways for climate change mitigation, it has not yet been deployed on the scale understood to be required, owing to a variety of technical, economic and commercial challenges.7
Carbon Capture and Storage (CCS)
CCS will be so important in our challenge against climate change because it has major advantages compared to the renewable energies available nowadays. Firstly, there is the possibility of retrofitting
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already existing power plants or industries at a moderate cost. Furthermore, the CCS landscape is rich of different methods and technologies (Figure 3), so there are different solutions to different situations and a certain extent of customization opportunity. The choice of technology is dependent on the requirements for product purity and on the conditions of the gas stream being treated (such as its temperature, CO2 partial pressure, concentration and the type and level of trace species or impurities). The major drawbacks of CCS are for sure of economic nature; especially in the first period massive investments will be needed to build a network of infrastructures which are essentially non-existent today. These infrastructures will have to encompass, beside the absorption plants, also transportation networks and storage facilities. Moreover, not all the methods and strategies that will be described herein are at their technological maturity; on the contrary, very few of them have reached the commercial level. In summary, although there has been a raising level of interest and research in the past years we are far from a thorough knowledge of the field neither on the economic nor on the technical side (without mentioning the political one).
There are four main approaches for carbon capture which all aim to produce a stream of pure CO2 that can be permanently stored or sequestered:
Pre-combustion capture:8 Coal or natural gas is subjected to a process of steam-reforming to produce syngas (H2+CO) which is cleaned from impurities and converted in a two-stage water gas shift reactor in H2+CO2. At this pressure and concentration, the CO2 is removable by physical adsorption and the H2 stream is sent to the burner.
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Post-combustion capture:9 Flue gases are collected and scrubbed in
an absorbent solution (usually amines), adsorbed with solid sorbents or separated with semipermeable membranes.
Oxy-combustion capture:10 Pure oxygen rather than air is used for combustion. This eliminates the large amount of nitrogen in the flue gas stream. After the particulate matter (fly ash) is removed, the flue gas consists only of water vapor and CO2 plus smaller amounts of pollutants such as sulphur dioxide (SO2) and nitrogen oxides (NOx). The water vapor is removed by cooling and compressing the flue gas. Additional removal of air pollutants leaves a nearly-pure CO2 stream that can be sent directly to storage.
Direct-Air capture:11 Large volumes of air are pushed in a sorbent system. Once the system has reached saturation, it is regenerated and the cycle starts again, as for the other types of capture. It is a difficult process due to the extreme low concentration of CO2 in the atmosphere.
Among these approaches only the post-combustion capture has reached commercial distribution being the oldest patented technology. The energy requirements of current CO2 capture systems are roughly ten to a hundred times greater than those of other environmental control systems employed at a modern electric power plant. This energy “penalty” lowers the overall (net) plant efficiency and significantly increases the net cost of CO2 capture.12 In general, the higher the power plant efficiency, the smaller are the energy penalty and associated impacts. For this reason, replacing or repowering an old, inefficient plant with a new, more efficient unit with CO2 capture can still yield a net efficiency gain that decreases all plant emissions and resource consumption. Thus, the net impact of the CO2 capture energy penalty
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must be assessed in the context of a particular situation or strategy for reducing CO2 emissions. As depicted in Figure 3 there are many types of carbon capture: chemisorption, adsorption, physisorption, carbon fixation and others.
Chemisorption
This is the process in which the CO2 is included in the bulk of a liquid by reacting with some specie present in solution, yielding new compounds (e.g. carbamate, bicarbonate, …).
The typical chemical absorbent for CO2 is a 20-30 wt% aqueous monoethanolamine (MEA) solution. It is particularly suited to low CO2 partial pressure applications and, consequently, has become the benchmark amine for CO2 capture from electricity generation; new sorbents are typically compared to it in terms of performances. MEA has good rates of CO2 mass transfer, is low cost and readily biodegradable but suffers from moderate rates of oxidative and thermal degradation and moderate levels of toxicity.13 It is also corrosive when used at higher concentrations and the overall process is energy demanding. Other amines systems have been assessed for their CO2 capture performances. Primary and secondary amines react with CO2 to form a carbamate or bicarbonate reaction product depending on the steric hindrance of the amine and the water concentration.14 Tertiary amines are forced to produce bicarbonate.
Other important solutions adopted for chemical absorption are aqueous solution of carbonates or alkaline media to sequestered CO2 in the form of bicarbonates.15 The most developed process based on this kind of
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compounds is depicted in Figure 2.16 This process comprises two chemical loops. The first loop captures CO2 using an aqueous solution of potassium hydroxide (KOH) forming potassium carbonate (K2CO3). In the second loop, called calcium loop (CaL), the carbonate (CO32-) is precipitated by reaction with calcium hydroxide (Ca(OH)2) to form calcium carbonate (CaCO3), restoring the initial solution of potassium hydroxide. The calcium carbonate is then calcinated (at 800-900°C) to liberate CO2, producing calcium oxide (CaO), which is hydrated to produce back the calcium hydroxide (Ca(OH)2).
Figure 2. Scheme of an absorption process with alkaline solution coupled with a calcium loop (CaL) process of desorption.
Newer systems are Ionic liquids (ILs) and Deep Eutectic Solvents (DES) which due to some interesting properties as low volatility and a wide range of functionalizations have raised a large interest in the past years and they can show remarkable capture performances.17,18 An interesting example could be a polyamines/ethylene glycol-based systems which show absorption of 0.93 molCO2/molamine (17-22% w/w), complete desorption with low temperature heating (110°C) and high thermal stability.19 These types of system will be further discussed in more depth.
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Adsorption
In these systems, CO2 is bound to a solid sorbent by weak interactions; this has the advantage of low energies of adsorption and desorption that translate in low cost of operation, reversibility of the process for a large number of cycles and usually great performances in terms of wt% of CO2 captured. The main adsorbents available are zeolites, MOFs (Metal-Organic Frameworks) and mesoporous silicas, such as MCM (Mobil Composition of Matter).20 All these technologies are still in a developmental stage and they have the drawback of being very expensive materials. They will have to show incredible performances to be adopted on large scales.
Physisorption
Also called physical adsorption, it is the process in which the CO2 is kept by a solvent via Van der Waals forces which barely perturb its structure. It is only feasible with high concentration CO2 streams (see pre-combustion capture) and it is performed in a wide range of pressure (from 3 to 70 atm). Several patents have been registered like Selexol,21 a solution of dimethyl ethers and polyethylene glycol; and Rectisol,22 a methanol solution kept a low temperature such as -40°C.
Carbon Fixation
These strategies come from disciplines outside chemistry as biology and geology. The primary proposals are forestation/afforestation, ocean fertilisation and mineral carbonation. Except for the first one, which is
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globally employed, the others are only proposals and need much more research to be adopted.
Methods of regeneration
The most common method of regeneration is the Temperature Swing Absorption (it is the one utilized to regenerate the MEA solutions). After absorption, the solution is heated, and the CO2 released is collected and sent to storage or utilized. Most of the chemical absorption systems adopt this method of regeneration. For solid sorbents, Pressure Swing Adsorption or Vacuum Swing Adsorption can also be employed, where the CO2 is released upon a drop of pressure in the sorbent chamber. At last, there is the pH Swing Absorption. It is an electrochemical method that exploit the different solubility of CO2 between alkaline and acid solutions. This method is still not largely utilized but is under intensive study.23,24
Storage and reuse
The most advanced method of storage of CO2 is the geological type, where supercritical CO2 is pumped in stable geological reservoirs as dismissed mines, salt domes and depleted oil deposits. For the latter one the technology is called EOR (Enhanced Oil Recovery) in which the filling of the deposit with CO2 enables the recovery of last traces of oil present. This kind of storage is not without its problem; for instance, a
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major concern is the possibility of massive leaking and even higher risk of earthquakes.
Another method under discussion is ocean storage in which the enormous pressures at the ocean floor would be exploited to maintain a layer of liquid CO2. This has raised a lot of concerns and its feasibility is not yet sure.
Obviously, the ideal solution would be to not have the need of storing the CO2, and because of this reason a lot of research has been devoted to its reuse in many different fields. However, this scenario is not entirely feasible since the amount of CO2 that should be captured vastly exceeds the current demand for carbon based products.25,26
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Pre-Combustion Post-Combustion Oxy-Combustion Direct Air Capture (DAC)
Types of capture
Methods of capture
Absorption 1)Amine solutions 2)Ammonia solutions 3)CaL 4)Carbonate solutions 5)Alkaline media 6)Ionic liquids7)Deep EutecticSolvents 8)Physisorption Adsorption 1)Zeolites 2)MOF 3)MCM Membranes 1)gas separation 2)gasabsorption Carbon fixation 1)Forestation 2)Ocean fertilization 3)Mineral carbonation
Geological storage Ocean storage EOR (enhanced oil recovery) Reuse
Storage
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Deep Eutectic Solvents and the Hole Model
DESs arose by the study of Ionic Liquids (ILs). While the latter are formed by one type of discrete cation and anion (as they can be considered low temperature molten salts), the former could have been described as a mixture of Lewis or Brønsted acids and bases which can contain a variety of anionic and/or cationic species.27 After the discovery that a mixture of choline chloride and urea showed a significant decrease of the melting point and great solvent properties,28 the term Deep Eutectic Solvent was coined. The definition changed to be a mixture of Donors and Acceptors, where the charge delocalization occurring through the weak interactions is responsible for the decrease in the melting point of the mixture relative to those of the individual components. Since then, several formulations have been investigated and nowadays there are four major classes of DESs that are summarized in Table 1.29
Type General Formula Terms
Type 1 𝐶𝑎𝑡+𝑋−𝑧𝑀𝐶𝑙𝑥 M= Zn, Sn, Fe, Al, Ga, In
Type 2 𝐶𝑎𝑡+𝑋−𝑧𝑀𝐶𝑙
𝑥∙ 𝑦𝐻2𝑂 M= Cr, Co, Cu, Ni, Fe
Type 3 𝐶𝑎𝑡+𝑋−𝑧𝑅𝑍 Z=CONH2, COOH, OH Type 4 𝑀𝐶𝑙𝑥+ 𝑅𝑍 = 𝑀𝐶𝑙𝑥−1+ ∙ 𝑅𝑍 + 𝑀𝐶𝑙𝑥+1− M=al, Zn and Z= CONH2,
OH
Table 1. Current classification of Deep Eutectic Solvents
Although this classification has shown to be comprehensive of most kind of DESs, others are still been proposed30 and considering the relative novelty of these compounds a proper, conclusive definition and, in consequence, a straight method of characterization, are still lacking. Besides these definitions, most DESs have in common several chemical and physical properties such as low toxicity, low vapor pressure, high
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viscosity, moisture stability and strong solvents properties towards different compounds. To better comprehend the nature of these systems many models have been explored but the one that showed the best agreement between theoretical calculations and experimental results is the Hole Model.
The Hole Model
Initially introduced to describe the properties of molten salts,31 the Hole Model has been successfully utilized to describe both ILs and DESs. This model assumes that ionic materials contain empty spaces that arise from thermally generated fluctuations in local density.32 The size of these empty spaces (holes) and especially the ratio between their size and the ions size dictate the different properties and behaviours of the system. For example, due to their thermal dependencethe average size of the holes in molten salts (high temperatures) is similar to that of the corresponding ion, so it is relatively easy for a small ion to move into a vacant site, and accordingly, viscosity of the liquid is low. However, the average size of holes will be smaller in lower temperature systems, which, coupled with a larger ion size, makes ion mobility difficult and explains why viscosities are higher in systems such as DESs. Following this reasoning, the model suggests an Arrhenius-type behaviour of viscosity as a function of temperature:
𝑙𝑛 𝜂 = 𝑙𝑛 𝜂
0+
𝐸𝜂𝑅𝑇 (1)
where 𝜂0 is a constant and 𝐸𝜂 is the energy for activation of viscous flow33. Since viscosity is limited by the low concentration of suitably
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sized voids, then it is logical to consider charge transfer in the same manner, hence, ion motion depends on the migration of holes in the opposite direction and conductivity is described in a similar fashion:
𝑙𝑛 𝜎 = 𝑙𝑛 𝜎
0−
𝐸𝜎𝑅𝑇 (2) Where 𝜎0 is a constant and 𝐸𝜎 is the energy for activation of ionic flow. An important consequence of this model is that the holes, rather than the ions, are responsible for charge transport.34 For this reason, the Walden rule is applicable for these systems. According to the Walden rule, the product of the molar conductivity and the fluidity (inverse of the viscosity) of a liquid is a constant:
Λ𝜂 = 𝑘
(3)Or in the more common logarithmic form:
𝑙𝑛 Λ = ln 𝑘 + ln(𝜂
−1)
(4)where Λ is the molar conductivity and k is a temperature dependent constant.35 The Walden rule was originally based on observations of the properties of dilute aqueous solutions but has been found to be applicable in ILs and DESs as well since they can be considered “holes solution” at infinite dilution.36 This empirical relation and the deviations from it can provide useful information about the ionicity and vapor pressure of a system. In a Walden plot (Figure 4), a diluted solution of potassium chloride (KCl) is usually taken for comparison because, due to almost identical sizes of cation and anion, it is described by the bisector of the first quadrant in the logarithmic plot.36 A system that shows a deviation of the slope is composed by a mixture of cations and anions of different sizes. Systems for which the plot lies over the
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bisector are referred to as “super-ionic”, meaning there is a lower degree of correlation between the motion of cation an anion, translating in higher conductivities. On the contrary, systems lying under the bisector are referred to as “poor-ionic” showing a higher degree of correlation between cation and anion, hence lower conductivities. The horizontal distance from the bisector simply indicates the fluidity (or viscosity) of the system.
Figure 4. Walden plot
Vapor pressure, as well as conductivity, is dependent on the formation of an “ideal” quasi-lattice. In the ideal quasi-lattice, the distribution of positive charge around negative charge gives the system a Madelung energy that is comparable to that of the ideal ionic crystal (i.e. KCl). The more ideal the quasi-lattice formed, the larger the Madelung energy of the liquid, hence the larger the energy that must be provided to extract
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an ion pair into the vapor state. Thus, ionic liquids lying on the ideal line will have the lowest vapor pressures.37 This model has shown to be very useful in the characterization of DESs and in suggesting strategies to improve some of their properties such as viscosity, conductivity and solvent capacity.38,39 More properties can be deduced from this model (e.g. diffusivity constants, surface tension, refractive indices), but since they have not been used to characterize the systems of interest they will not be discussed here.
Even though the research focused extensively on DESs, the two main systems studied do not fit in the current classification and, as it will be shown in the discussion, the characterization of these two systems and a clear definition is not trivial.
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Process of discovery
The research carried out here started by replicating a system from a previous thesis work;40 a DES (Deep Eutectic Solvent) composed of choline chloride/urea/lysine/ethylene glycol (1:2:1:3)(throughout the text these proportions are intended as molar ratios). This system absorbs CO2 producing carbamate due to the nucleophilic attack of the amine group of the amino acid to the carbon atom of carbon dioxide. This formulation shows high degree of absorption with the downside of being highly viscous and having the tendency to form a precipitate during carbonation.
The protonated form of the amines of the amino acid builds a network of hydrogen bonds, causing the high viscosity of the system. The deprotonated form of these sites should instead decrease the number and the strength of those interactions. In literature, the most common ways of deprotonation use amines, polyamines19 and superbases.41 Triethylamine, tributylamine and tetramethylammonium chloride have been screened. These formulations form biphasic systems which are troublesome to treat and the volatility of these amines does not allow heating.
Therefore, the focus has shifted to other bases. Wishing to utilize cheap and low toxic compounds potassium hydroxide (KOH) seemed a proper candidate. Simple systems like KOH/Lysine (1:1) form liquids (DES) with very high viscosity that has been hard to lower even adding ethylene glycol (EG). After many unsuccessful attempts to lower the viscosity, the amino acid species were removed from the formulations, obtaining a very simple system of KOH/EG. Even though this system
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is already a liquid at the molar ratio 1:1, a formulation with three equivalents of ethylene glycol to one of potassium hydroxide has been observed to be a good compromise between viscosity, absorption performance and air stability (see discussion).
The change in methodology led to a research in carbon capture by carbonate and bicarbonate formation.42 Actually, also in amino acid-based or amines-acid-based solvents, carbonates formation is competitive to the carbamate formation depending on the water content of the system43 and carbamates are usually more of a transient species in absorption.14,44 Indeed, the exclusive carbamate formation is observed only under complete anhydrous conditions.45
Several attempts have been conducted to further decrease the viscosity of the system, from adding dimethyl sulfoxide (DMSO),46 to varying the cation dimension using sodium hydroxide (NaOH) and caesium hydroxide (CsOH). With NaOH the system is highly viscous and very unstable to air exposure and although the viscosity has been reduced with both the other methods the advantages have not appeared to outweigh the downsides. The addition of DMSO does decrease the viscosity but it does not significantly enhance the absorption capacity and decreases the solubility of carbonates produced. In fact, it could have brought additional complication, adding a process of filtration to the regeneration cycle. Similarly, the use of caesium hydroxide decreased the viscosity of the system, but it has not shown great absorption improvements in comparison to the potassium hydroxide. Anyway, it was discarded because its scarce natural abundance and, consequently, its high cost. The will to avoid the use of water is dictated by the fact that one of the major drawbacks of the current methods of
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carbon capture (i.e. monoethanolamine aqueous solution) is the water high heat capacity (cp = 4179.6 J/K*kg at 25°C), which makes the process energetically demanding. Therefore, solvents with lower heat capacity (e.g. ethylene glycol, cp = 2408 J/K*kg at 25°C47) are attractive possible replacements.
Afterward, the possible mechanisms of desorption and regeneration have been investigated. Carbonates have much higher chemical stability then carbamates, though, and the usual methods of desorption as temperature increase or vacuum pumping have shown to be ineffective. Anyway, there is another desorption method which has received increasing attention in recent years: the pH swing.
The pH swing consists in lowering the pH of the sorbent by adding acids or via electrochemical ways;23,24,48 in so doing, the carbonates are protonated to carbonic acid, which decomposes, enabling the release of the captured CO2. There are many organic and inorganic acids that could be in line with a cheap and green system for carbon capture. After few attempts, formic acid and oxalic acid seemed to be the ideal candidates. Not only because they can convert all the carbonates back to CO2, but they themselves could be synthesized from CO2. Specifically, it has been shown that electroreduction of CO2 can promote the production of formic and/or oxalic acid.49,50 The addition of these acids to our carbonated system allowed the quantitative release of CO2 and forms a white precipitate of potassium formate or potassium oxalate. At this point a method of regeneration was needed. The oxidation of these compounds back to CO2 seemed to be the most logic follow-up, theoretically generating a closed loop of reduction/oxidation for system regeneration. After a screening of different oxidizing agents,
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we thought that a reaction of hydrogen peroxide found in literature could be useful. The reaction involves a superoxide (OH2-) moiety generated in mild alkaline condition that generates hydroxyl ions as product. This reaction was really promising because it could remove the precipitate from the system and regenerate the alkaline conditions at the same time. Unfortunately, under our experimental conditions the auto-oxidation of hydrogen peroxide prevailed and the conversion of oxalate into CO2 was negligible. Therefore, electrochemical methods to regenerate the original sorbent have been investigated (see results and discussion).
Among others, boric acid has also been used to regenerate the system. This acid enables the complete release of CO2 too, but it does not form any precipitate. Out of curiosity, after the addition of boric acid the system was carbonated again and although in much less extent the system reabsorbed CO2. After trying different ratios and different water contents the best formulation seemed to be KOH/H3BO3/EG (1:1:3). This system absorbs way less than the previous system, but it does not form precipitates, has low viscosity and it is stable if exposed to air. Besides, the most promising feature of this system has been its capacity of desorption and regeneration (see results and discussion).
In the final analysis, the two systems that became the candidates to further investigation and characterization are KOH/EG (1:3) and KOH/H3BO3/EG (1:1:3)
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Experimental section
Synthesis and carbonation
All the KOH/EG liquid systems were prepared by mixing the components under a N2 atmosphere to prevent the mixture to turn yellow (see discussion). The KOH/H3BO3/EG systems were prepared by adding the needed amount of boric acid to the freshly prepared KOH/EG liquid; this mixture is stable in air and no N2 atmosphere was needed.
Carbonation was conducted, under constant magnetic stirring, by connecting the mixture vessel, after vacuum was achieved, to another vessel where sulphuric acid was dripped onto sodium carbonate to release CO2 until a pressure of 1 atm was achieved. Due to difficulties in maintaining a constant CO2 pressure and a scarce reproducibility, this method was abandoned, and the mixture vessel was then connected directly to a balloon of pure CO2. Gravimetric measurements were taken to evaluate the amount of CO2 absorbed (technical balance, ±0.01 g).
NMR
NMR spectra were obtained on a Bruker Avance II DRX400 instrument equipped with a BBFO broadband probe. Chemical shifts (expressed in parts per million) are referenced to external standard (HDO in D2O). In
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most of the case, no lock and no spinning were used to analyse the pure mixtures.
Electrochemical analysis
Cyclic voltammetry was performed in a 0.05 M sodium sulphate (Na2SO4) solution using a glassy carbon working electrode, a platinum counter-electrode and a standard hydrogen electrode (SHE, E0=0.00 V) as reference.
For chronoamperometry and chronopotentiometry the reference was substituted with a silver chloride electrode (Ag/AgCl, E0=0.141 V)
Thermogravimetric Analysis
Thermograms of the mixtures were obtained with a Thermogravimetric analyser Q5000 V3. The samples were heated from 25°C to 240°C with IR lamps under a pure N2 flow.
Conductivity
Conductivities were measured from 25°C to 75°C using an HACH sensIONTM+ MM374 conductivity meter with platinum electrode and a LAUDA E100 thermostat.
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Viscosity
Viscosities were measured from 25°C to 75°C with a Fungilab EXPERT rotational viscosimeter connected to a LAUDA E100 thermostat.
Density
Densities were measured to compare the behaviour of the different systems on a Walden plot. The density values at different temperatures (from 25°C to 75°C) were obtained by measuring the weight of the sample in a 2.00 mL volumetric flask after thermal equilibrium was reached using a LAUDA E100 thermostat.
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Results
Ethylene Glycol
1H NMR spectrum (Figure 5), 13C NMR spectrum (Figure 6) and Thermogravimetric analysis (Figure 7) of pure ethylene glycol were obtained as benchmark. Water content was also established by measuring the weight loss after drying the ethylene glycol in stove at 90°C for 24 hours. The result was 3.18 % w/w.
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Figure 6. 13C NMR spectrum of pure ethylene glycol
Figure 7. TGA of pure ethylene glycol
The two peaks in the 1H NMR spectrum of pure ethylene glycol (Figure 5) have been assigned to the hydrogen nuclei belonging to the hydroxyl
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groups (5.33 ppm) and to the methylene groups (3.75 ppm); this is suggested not only by the values of chemical shift but also by the relative intensity of the peaks being the hydroxyl peak about half of the other, perfectly reflecting the proportion of the chemical formula. the 13C NMR spectrum shows only one peak at 63.21 ppm, coherently with the symmetry of the molecule. The thermogram shows two peaks in the rate of weight loss at 64.60°C and 131.75°C leading to complete degradation with no residue.
KOH/EG
This system was characterized by measuring its physical properties such as density (Figure 8), viscosity (Figure 9) and conductivity (Figure 10). The chemical composition was instead investigated by the acquisition of 1H NMR spectrum (Figure 11), 13C NMR spectrum (Figure 12) and thermogravimetric analysis (Figure 13).
Figure 8. Density plot vs temperature of potassium hydroxide and ethylene glycol (1:3) 1.225 1.23 1.235 1.24 1.245 1.25 1.255 1.26 20 30 40 50 60 70 D e n si ty (g/m L) Temperature (°C)
28
Figure 9. Viscosity plots vs temperature (normal scale on the left, logarithmic scale on the right) of potassium hydroxide and ethylene glycol (1:3)
Figure 10. Conductivity plots vs temperature (normal scale on the left, logarithmic scale on the right) of potassium hydroxide and ethylene glycol (1:1:3)
Figure 11 1H NMR spectrum of potassium hydroxide and ethylene glycol (1:3)
0 50 100 150 200 250 300 350 23 33 43 53 63 73 Vi sco si ty (cP ) Temperature (°C) 3 3.5 4 4.5 5 5.5 6 0.00285 0.00305 0.00325 ln (η) ( cP) 1/T (K-1) 2 4 6 8 10 12 14 16 18 15.00 25.00 35.00 45.00 55.00 65.00 75.00 Co n d u ctiv ity (m S/ cm ) Temperature (°C) 1.3 1.5 1.7 1.9 2.1 2.3 2.5 2.7 2.9 0.00286 0.00306 0.00326 ln (CE ) m S/ cm 1/T (K-1)
29
Figure 12. 13C NMR spectrum of potassium hydroxide and ethylene glycol (1:3)
Figure 13. TGA of potassium hydroxide and ethylene glycol (1:3)
In the 1H NMR spectrum of KOH/EG (Figure 11), the peak of methylene hydrogens shows a minor shift to lower value of 3.48 ppm. Since the spectra have been obtained without the use of any solvent to
30
perform the lock this type of shift is not considered relevant. The hydroxyl peak instead shows a non-negligible shift with a value of 6.65 ppm. The 13C NMR spectrum shows only one peak at 63.36 ppm. The thermogram shows two major peaks of weight loss rate at 58.17°C and 130.24°C. Two minor ones, and less resolved, can be seen at 89.42°C and 221.61°C. A solid residue remained after heating, likely of potassium glycolate.
KOH/H
3BO
3/EG
The same analysis as for the KOH/EG system were conducted for this mixture (except the 1H NMR spectrum): density (Figure 14), viscosity (Figure 15), conductivity (Figure 16), 13C NMR spectrum (Figure 17) and TGA (Figure 19). The presence of boric acid allowed a further investigation through the acquisition of 11B NMR spectrum (Figure 18).
Figure 14. Density plot vs temperature of potassium hydroxide, boric acid and ethylene glycol (1:1:3) 1.27 1.28 1.29 1.3 1.31 1.32 20 30 40 50 60 70 D e n si ty (g/m L) Temperature (°C)
31
Figure 15. Viscosity plots vs temperature (normal scale on the left, logarithmic scale on the right) of potassium hydroxide, boric acid and ethylene glycol (1:1:3)
Figure 16. Conductivity plots vs temperature (normal scale on the left, logarithmic scale on the right) of potassium hydroxide, boric acid and ethylene glycol (1:1:3)
10 20 30 40 50 60 70 80 90 100 110 23 33 43 53 63 73 Vi sco si ty (cP ) Temperature (°C) 2.5 3 3.5 4 4.5 5 0.0029 0.0031 0.0033 ln (η) ( cP ) 1/T (K-1) 2 3 4 5 6 7 8 9 10 15.00 25.00 35.00 45.00 55.00 65.00 75.00 Co n d u ctiv ity (m S/ cm ) Temperature (°C) 1 1.2 1.4 1.6 1.8 2 2.2 2.4 0.00286 0.00306 0.00326 ln (CE ) m S/ cm 1/T (K-1)
32
Figure 17. 13C NMR spectrum of potassium hydroxide, boric acid and ethylene
glycol (1:1:3)
Figure 18. 11B NMR spectrum of potassium hydroxide, boric acid and ethylene
33
Figure 19. TGA of potassium hydroxide, boric acid and ethylene glycol (1:1:3)
The 13C NMR spectrum still shows only one peak at 62.71 ppm while the 11B NMR spectrum shows two peaks at 9.63 ppm and 6.20 ppm. The thermogram reveals a significant weight loss until about 110°C with a maximum rate at 54.81°C. A second weight loss is observed from 150°C. The jagged band of the weight loss rate should suggest the formation of bubbles and spurts that can alter the weight measurement.
KOH/EG + CO
2The presence of a significant quantity of precipitate in this system has prevented the measurement of the physical properties as for the other systems. In particular, when heated, this system turns into a gel-like compound with very high viscosity leading to radical change in behaviour of viscous flow (non-newtonian fluid) and ion mobility. Only
34
density is reported (Figure 20) to show the change occurring above 40°C. Also the 13C NMR spectrum (Figure 21) and the TGA (Figure 23) were obtained.
Figure 20. Density plot vs temperature of carbonated potassium hydroxide and ethylene glycol (1:3)
Figure 21. 13C NMR spectrum of carbonated potassium hydroxide and ethylene
glycol (1:3) 1.23 1.24 1.25 1.26 1.27 1.28 1.29 1.3 20 30 40 50 60 70 D e n si ty (g/m L) Temperature (°C)
35
Figure 22. 13C NMR spectrum of potassium bicarbonate and ethylene glycol
36
The 13C NMR spectrum shows five peaks at 60.08 ppm, 62.67 ppm, 66.57 ppm, 158.24 ppm and 167.73 ppm. The last peak can be likely assigned to the bicarbonate species, as confirmed by the 13C NMR spectrum of a potassium bicarbonate (KHCO3) in ethylene glycol (Figure 22, see discussion). The thermogram shows an almost steady weight loss until about 140°C with three poorly resolved peaks of weight loss rate at 52.10°C, 86.48°C and 128.39°C.
KOH/H
3BO
3/EG + CO
2This system was characterized by all the analyses conducted before. Physical properties: density (Figure 24), viscosity (Figure 25), conductivity (Figure 26). 13C NMR spectrum (Figure 27), 11B NMR spectrum (Figure 28) and TGA (Figure 29).
Figure 24. Density plot vs temperature of carbonated potassium hydroxide, boric acid and ethylene glycol (1:1:3)
1.26 1.27 1.28 1.29 1.3 1.31 1.32 1.33 20 30 40 50 60 70 D e n si ty (g/m L) Temperature (°C)
37
Figure 25. Viscosity plots vs temperature (normal scale on the left, logarithmic scale on the right) of carbonated potassium hydroxide, boric acid and ethylene glycol (1:1:3)
Figure 26. Conductivity plots vs temperature (normal scale on the left, logarithmic scale on the right) of carbonated potassium hydroxide, boric acid and ethylene glycol (1:1:3) 10 20 30 40 50 60 70 80 90 100 110 24 44 64 Vi sco si ty (c P) Temperature (°C) 2.6 3.1 3.6 4.1 4.6 0.0029 0.0031 0.0033 ln (η) ( cP) 1/T (K-1) 2 3 4 5 6 7 8 9 10 15.00 25.00 35.00 45.00 55.00 65.00 75.00 Co n d u ctiv ity (m S/ cm ) Temperatura (°C) 1 1.2 1.4 1.6 1.8 2 2.2 2.4 0.00286 0.00306 0.00326 ln (CE ) m S/ cm 1/T (K-1)
38
Figure 27. 13C NMR spectrum of carbonated potassium hydroxide, boric acid and
ethylene glycol (1:1:3)
Figure 28. 11B NMR spectrum of carbonated potassium hydroxide, boric acid and
39
Figure 29. TGA of carbonated potassium hydroxide, boric acid and ethylene glycol (1:1:3)
13C NMR spectrum shows five peaks at 60.38 ppm, 62.71 ppm, 66.70 ppm, 158.71 ppm and 160.33 ppm while the 11B NMR spectrum consists in three peaks at 0.36 ppm, 6.31 ppm and 9.55 ppm. The thermogram shows a significant weight loss until about 120°C with the maximum rate of weight loss at 55.79°C. Above 140°C another jagged profile of the rate of weight loss is observed as in the system before carbonation.
The desorption mechanism of this system was investigated through 13C NMR spectrum (Figure 30), 11B NMR spectrum (Figure 31) and gravimetric analysis (Figure 29).
40
Figure 30. 13C NMR spectrum of carbonated potassium hydroxide, boric acid and
ethylene glycol (1:1:3) after 30 minutes of heating at 60-70 °C under vacuum
Figure 31. 11B NMR spectrum of carbonated potassium hydroxide, boric acid and
41
The 13C NMR spectrum shows two very close peaks at 62.08 ppm and 62.48 ppm, likely due to a poor shimming and a relevant viscosity. The 11B NMR spectrum shows two peaks at 5.88 ppm and 9.36 ppm.
Figure 32. 11B NMR spectrum of dehydrated KOH:H3BO3:EG (1:1:1) dissolved in
EG
42
Figure 32 shows the 11B NMR spectrum of KOH:H3BO3:EG (1:1:1) which has been dehydrated by heating the mixture under vacuum until a soft white solid was formed. The spectrum has been obtained by dissolving the solid in a minimal quantity of ethylene glycol. The spectrum shows only one peak at 9.49 ppm. Figure 33 shows the 11B NMR spectrum of K[B(OH)4] in D2O. This compound was obtained by the crystallization of an aqueous solution in which equal quantities of KOH and H3BO3 were previously mixed. The spectrum shows only one peak at 6.16 ppm although it seems slightly contaminated by other species. These spectra have been used to investigate the assignment of the peaks of the system KOH:H3BO3:EG (1:1:3).
43
Electrochemical results
In order to explore the pH swing with oxalic acid and the consequent electrolysis of the product as a desorption step, the electrochemical analysis of potassium oxalate (K2C2O4) in water is reported below.
Figure 34. Cyclic voltammetry of a sodium sulphate solution 0.05 M (Blank);
Reference electrode: Standard hydrogen electrode (SHE) (E0=0.000 V) activity of
H+=1 Molar -10.00 -5.00 0.00 5.00 10.00 15.00 20.00 -1.000 -0.500 0.000 0.500 1.000 1.500 Cu rr e n t/ µ A Potential/V
44
Figure 35. Cyclic voltammetry of potassium oxalate (0.001M)/ sodium sulphate (0.05 M) solution; Reference electrode: Standard hydrogen electrode (SHE)
(E0=0.000 V) activity of H+=1 Molar
Figure 36. Chronoamperometry (electrolysis) of potassium oxalate (0.001M)/
sodium sulphate (0.05 M) solution; Reference electrode: Ag/AgCl (E0=0.141 V)
-5.00 0.00 5.00 10.00 15.00 20.00 25.00 0.000 0.500 1.000 1.500 Cu rr e n t (µ A ) Potential (V) -1,000 0 1,000 2,000 3,000 4,000 5,000 6,000 7,000 0.000 500.000 1000.000 1500.000 2000.000 Cu rr e n t (µ A ) Time (s)
45
Figure 37. Chronopotentiometry of potassium oxalate (0.001M)/ sodium sulphate
(0.05 M) solution; Reference electrode: Ag/AgCl (E0=0.141 V)
The cyclic voltammetry (Figure 35) was performed to estimate the electric potential needed to oxidate the oxalate into CO2. The analysis indicated a value of E0=1.325 V. Afterward, an electrolysis or chronoamperometry (Figure 36) was conducted to test the feasibility of the process. The applied voltage was initially set to the value suggested by the cyclic voltammetry. For unknown reason, after about three minutes random spikes of current (negative and positive) were observed and the voltage had to be decreased to about 0.86 V referred to the silver chloride electrode ( ̴ 1.001 V referred to SHE) to have a smooth decrease of current. To test this value, a chronopotentiometry was conducted (Figure 37). A current of 4.0 mA was set for one hour. After few seconds, the voltage settled at about 0.85 V referred to the silver chloride electrode ( ̴ 0.991 V referred to SHE). Being interested just to the proof of concept of this process, no further investigation was conducted. The reaction has also been monitored by measuring the pH
0.00 0.10 0.20 0.30 0.40 0.50 0.60 0.70 0.80 0.90 0 1000 2000 3000 Pot e n tial ( V) Time (s)
46
of the solution at the counter electrode (where hydroxide anions should be forming) before and after electrolysis. The solution of sodium sulphate had an initial pH of 5 and after 30 minutes it raised to a value of 12. The big increase was also due to the small volume of the solution but nevertheless it qualitatively proved the success of the reaction.
Walden Plot
For all the systems except the carbonated KOH/EG the logarithm of molar conductivity versus the logarithm of fluidity was plotted (Figure 38) obtaining useful information about the internal structure and ionicity of the system.
Figure 38. Walden Plot
-2 -1.5 -1 -0.5 0 0.5 1 1.5 2 -2 -1 0 1 2 ln (Λ ) (S *c m 2/m o l) ln(1/η) (P-1) KOH:B(OH)3:EG (1:1:3) KOH:EG (1:3) KOH:B(OH)3:EG (1:1:3)+CO2
47
Melting Points
Table 2 shows the different melting points of the systems before and after carbonation, along with the melting points of the single compounds. The melting points of the different mixtures have been measured by freezing the systems in a bath of acetone and solid CO2 (-77°C) with a thermometer inside. The temperature of melting was taken when the thermometer could move inside the liquid. The results have been expressed rather as an interval than as a precise temperature for two reasons. Firstly, the method is not accurate, and many measurements have been taken to reduce the error. Secondly, these systems do not show a smooth transition from solid to liquid, but they have a plastic-like behaviour in the range of temperature showed in Table 2. System Tfus (°C) Ethylene glycol -1351 Potassium hydroxide 36052 Boric acid 170.953 KOH/EG (1:3) -40/-34 KOH/EG (1:3) + CO2 -55/-48 KOH/H3BO3/EG (1:1:3) -53/-46 KOH/H3BO3/EG (1:1:3) + CO2 -52/-48
48
Discussion
The significant shift of the hydroxyl peak from the spectrum of pure ethylene glycol to the spectrum of the KOH/EG suggests a fast equilibrium of deprotonation, as the peak is the weighted mean of the peaks of free hydroxide anions and the hydroxyl groups of the ethylene glycol; this is also suggested by the increased intensity of the peak compared to pure ethylene glycol. The high rate of this equilibrium is validated by the fact that the 13C NMR spectra (Figure 6-Figure 12) of the two systems are basically identical with only one peak at 63.21 ppm in pure ethylene glycol and a peak at 63.36 ppm in the system with KOH. So, the system KOH/EG is expected to undergo, at least by a certain extent, the reaction:
Scheme 1:
The KOH has always been mixed in pellets, without previous grinding.
This slows the dissolution in EG but it reduces the water content introduced which, for this reason, is considered to be negligible. Therefore, in order to calculate how much ethylene glycol is deprotonated, the information about the water content in both the EG and in the final system is needed. Unfortunately, Karl-Fischer titration proved to be not suitable for such alkaline systems. The thermogravimetric analysis was performed mainly to investigate this aspect of the different systems, but it also led to inconclusive results. In fact, the sample of pure ethylene glycol (Figure 7) was kept at the air
49
for too long before the analysis, leading to an inconsistently high water content, even higher than the KOH/EG system. For this reason, the water content calculated by thermal dehydration of the ethylene glycol was then considered (3.18% w/w). Nevertheless, even using this value the thermogravimetric measurements did not produce valuable information about the water content. In the TGA of the system KOH/EG (Figure 13), considering the first weight loss as the water evaporation we can calculate the ratio of deprotonated to protonated glycol. The initial weight was 12.3320 mg. Considering the loss of the 18.91% of weight gives a water content of 2.3320 mg. Knowing the molar ratio of the components the amount of ethylene glycol can be calculated: Scheme 2: 𝑀 = 𝑚𝑜𝑙𝐾𝑂𝐻 ∙ 𝑃𝑀𝐾𝑂𝐻 + 𝑚𝑜𝑙𝐸𝐺 ∙ 𝑃𝑀𝐸𝐺 = =1 3∙ 𝑚𝑜𝑙𝐸𝐺 ∙ 𝑃𝑀𝐾𝑂𝐻 + 𝑚𝑜𝑙𝐸𝐺 ∙ 𝑃𝑀𝐸𝐺 𝑚𝑜𝑙𝐸𝐺 = 1 𝑀 3 ⁄ ∙ 𝑃𝑀𝐾𝑂𝐻 + 𝑃𝑀𝐸𝐺
Where M is the initial weight of the sample and PM are the molecular weights of the compounds. From the mass of ethylene glycol and its degree of hydration a content of water of 0.3014 mg is obtained suggesting the water produced by deprotonation to be 2.0306 mg (0.13 mmol). The theoretical maximum molar quantity of water produced is equal to the moles of KOH (0.059 mmol), obtainable by a similar calculation to the one showed in Scheme 2. The comparison of these two values suggests an extent of deprotonation of 222%. The evaporation of the water could have shifted the deprotonation
50
equilibrium completely toward the products, but this is not enough to explain this result; most likely, in such small amounts, a brief exposure to the atmospheric humidity is enough to significantly alter the data. It is also possible, especially at higher temperature, that the peak is “contaminated” by the loss of ethylene glycol which, in the condition of the N2 flow, shows a much greater tendency to evaporation compared to ambient conditions. This is proved by the TGA of this compound (Figure 7) that shows complete evaporation (with no residue) way before the boiling temperature reported in the literature (197.6°C).51 Unfortunately, no other method to investigate the water content to understand the degree of deprotonation without perturbing such equilibrium has been found.
As previously said, this system had to be prepared under a N2 atmosphere, in order to prevent a reaction that brought the mixture to become yellow at first and then dark yellow/brown. This reaction is believed to be caused by the atmospheric oxygen which could react with the ethylene glycol leading to some oxidation product. Although the effect of this reaction is clearly visible by eye, no difference in terms of NMR spectra or absorption capacity have been observed. The most likely explanation of this behaviour is that the extent of the reaction is very small, negligible indeed, but the products obtained have very high light absorption coefficient, giving the mixture an intense colour even at very low concentration. This phenomenon was not further investigated due to the absence of effects on the absorption capacity of the system and the relatively easy solution to it (N2 atmosphere). This system has a density around 1.25 g/cm3 (Figure 8) at room temperature
51
and a viscosity of 330 cP1 (Figure 9) which is very sensitive to a change in temperature showing an Arrhenius type dependence on it. The conductivity (Figure 10) shows a clear dependence from the temperature, going from 4.0 mS/cm at room temperature to 16.27 mS/cm at 75°C.
Besides its simple composition, low cost, and low toxicity, this system has been considered promising by its high absorption capacity which it will be expressed as molCO2/molKOH. For this system, the maximum value obtained is 0.90. The absorption mechanism of this system has been investigated primarily through the comparison of 13C NMR spectra. The spectrum of the carbonated system (Figure 21) shows other four peaks, the most intense of which at 62.67 ppm, similarly to that previously assigned to the methylene carbons of the ethylene glycol. Its intensity is likely due to the stoichiometric excess of EG with respect to the KOH, which is ultimately responsible of the absorption mechanism. The two peaks closer to the peak of ethylene glycol (60.08 ppm and 66.57 ppm) are believed to be proof of the mechanism of absorption showed in Scheme 3. The presence of the carbonate group breaks the symmetry of ethylene glycol and the carbon nuclei of the backbone are not equivalent anymore. To assign the two peaks at higher chemical shifts the 13C NMR spectrum of potassium bicarbonate in ethylene glycol (Figure 22) was obtained. This spectrum shows a peak at 167.88 ppm which is consistent to the peak observed in the 13C NMR spectrum of the carbonated KOH/EG system at 167.73 ppm, indicating a mechanism of absorption via bicarbonate formation (Scheme 4). The
1 As reference, the viscosity values of water, oil and honey are respectively 1 cP at 20°C, 50.2 cP at 26°C and 103-104 cP at 20-30°C.
52
other peak at 158.24 ppm has been assigned to the carbonated glycol, in particular to the carbon originally belonging to the absorbed CO2.
Scheme 3:
Scheme 4:
The system precipitates during carbonation. After several attempts of regeneration via heating and vacuum, only the reduction of the peak at 158.24 ppm has been observed, but a complete disappearance of all the peaks has proved impossible. Therefore, the pH swing technique was explored. The details of this method of regeneration are outlined at the end of the discussion.
The second system, KOH/H3BO3/EG (1:1:3), quite cheap and “simple” in its formulation, is very complex in its internal structure. The addition of the boric acid also seems to establish rapid equilibria with the other species in the mixture. This can be implied by the 13C NMR spectrum of the system (Figure 17) that still shows only one peak at 62.71 ppm, consistent with the previous ones. This indicates the absence of stable compounds of boric acid with ethylene glycol. Indeed, without the presence of KOH the boric acid is practically insoluble in EG. The speciation of boric acid is extremely rich and complex54 (at least in
53
aqueous solution) and very sensitive to pH.55 The simplest reaction that can occur in the mixture is shown in Scheme 5.
Scheme 5:
Nevertheless, it proved not to be the only one since the 11B NMR spectrum (Figure 18) shows two broad and partially overlapping peaks. The assignment of these two peaks is not trivial, as references in ethylene glycol cannot be found in literature. One peak is almost for sure the borate of Scheme 5, but for the other a possibility could be: Scheme 6:
If this compound is indeed formed, it must be in rapid equilibrium with other species since the 13C NMR spectrum does show only one peak. In Scheme 7 is depicted the possible equilibrium of isomery between the specie in Scheme 6 and the Lewis adduct. Assuming the presence of this specie, it is then reasonable to assume that the glycolate is under a fast equilibrium with its protonated form resulting in only one peak in the 13C NMR spectrum:
54
Scheme 7:
This adduct could establish a fast equilibrium between the protonated and deprotonated glycol explaining the single peak in the 13C NMR spectrum. The TGA of this system (Figure 19) does provide some information that could validate this hypothesis.
Generally, most of the weight loss occurs at 130°C and it is thought to be due to ethylene glycol evaporation. In this system the most significant weight loss only starts at higher temperature (150°C), and this could be interpreted as a proof of a stronger interaction between the ethylene glycol and boron species, with a consequent increase in stability of the mixture to thermal degradation. The stronger association between the species is also suggested by the Walden plot of this system. The Walden points lie under the bisector, indicating a poor-ionic liquid, meaning a higher degree of association between the ions that can lead to higher reticular energy and lower vapor pressure. The physical analysis indicates density at room temperature of 1.31 g/cm3 (Figure 14) a viscosity around 110 cP (Figure 15), much lower than the previous system and a conductivity of 3.2 mS/cm (Figure 16). All these properties are consistent with a stronger ionic association and a higher water content.
More information about this system can be deduced by the study of the absorption and desorption mechanism. The system shows a much lower absorption with respect to the KOH/EG system: 0.15 molCO2/molKOH. No precipitation is observed during carbonation and the properties of
55
the system are not altered when compared to the KOH/EG/H3BO3, as it can be seen from the physical properties (Figure 24, Figure 25, Figure 26). Maybe the most interesting information come from the NMR spectra, both 13C and 11B. The 13C NMR spectrum (Figure 27) is very similar to the one of the carbonated KOH/EG, with the exception that the peak previously at 167.73 ppm has shifted to 160.61 ppm. The spectrum seems to suggest that the absorption still involves the carbonation of the deprotonated ethylene glycol (158.71 ppm) but not the formation of the simple bicarbonate. In fact, the peak at 160.61 ppm could be interpreted as the formation of a new species between the CO2 and some species of boron. The 11B NMR spectrum (Figure 28) instead still shows a peak at 9.55 ppm qualitatively unaltered while the peak at 6.31 ppm shows a massive decrease in intensity compared to the original system and the appearance of a third peak at 0.36 ppm. In the literature the product of the reaction:
Scheme 8:
is considered to be only a transient state 56,57 and it acts as a catalyst for the process of absorption.58,59 However, these considerations are valid in aqueous solutions and it has been observed that some species of boron, and especially the borates, can form stable compounds with diols.60 Therefore, it is possible that one of the peak could be assigned to the borate and the peak that appears at 0.36 ppm is indeed the
56
compound depicted in Scheme 8: which can be stabilized by the ethylene glycol.
The assignment of the peaks is based on the analysis of the dehydrated KOH:H3BO3:EG (1:1:1) in EG and K[B(OH)4] in D2O. For the former, it is reasonable to believe that with such molar ratios and after an intensive dehydration the extent of deprotonation of the EG should be almost complete. In such conditions, the formation of the Lewis adduct should be the most favoured. The spectrum of this system shows a peak at 9.49 ppm which is consistent with one of the peaks of the system KOH:H3BO3:EG (1:1:3). But this mixture also shows another peak at 6.20 ppm which is instead consistent with the spectrum of the potassium tetrahydroxyborate (K[B(OH)4]), indicate the presence of the borate anion, even though the solvent is different.
The TGA analysis of this mixture (Figure 29) does not show significant changes from the system before carbonation and it does not bring new information about it. The absence of the bicarbonate seems to be validated by the process of desorption. Indeed, when heating (at 70-90°C) and vacuum were applied on the carbonated system KOH/EG, the peak at 158 ppm (the carbonated glycol) seemed to decrease, and with intense treatment it eventually disappeared, while the peak at 167 ppm (the bicarbonate) never disappeared and the reduction in intensity was always negligible. In the carbonated KOH:H3BO3:EG system, the treatment with heat (60-70°C) and vacuum led to a complete desorption of CO2 and regeneration of the original mixture. In the 13C NMR spectrum after desorption (Figure 30) all the peaks previously assigned to carbonation phenomena have disappeared and only the peak associated to the ethylene glycol remains, even though it shows a
57
double peak. The very narrow distance of chemical shift between the two (0.4 ppm) could be indicative of an artifact of the measurement due to bad shimming. The 11B NMR spectrum (Figure 31) shows the complete disappearance of the peak associated to the carbonated species and the return to only the two peaks at the same chemical shift as in the original mixture. The only difference is the relative intensity of the peaks which is opposite to the spectrum before absorption. Besides, the absorption/desorption cycles did not reveal major changes in absorption, suggesting that no major structural transformations occur during the regeneration process.
Figure 39. CO2 Absorption-(CO2/H2O) Desorption cycles of the system
58
Initial weight (g) Final weight (g) Δweight (g) Absorption %
30.6076 30.7442 0.1366 14.52 30.4187 30.5551 0.1364 14.50 30.1229 30.2624 0.1395 14.83 30.0312 30.1004 0.0692 7.36 29.7982 29.8194 0.0212 2.25 29.4687 29.4738 0.0051 0.54 29.1305 29.1336 0.0031 0.33
Table 3. Absorption capacities as a function of water content. The highlighted line
corresponds to the initial formulation KOH:H3BO3:EG (1:1:3)
Figure 39 displays a series of absorption/desorption cycles at different water contents as they were conducted. Table 3. Absorption capacities as a function of water content instead reports the sample weight before the absorption (indicative of the water content, the higher the weight, the higher the water content), the weight of the sample after absorption, the weight increment (i.e. grams of CO2 absorbed) and the absorption capacity expressed as molCO2/molKOH. Specifically, it has been observed that at the original molar composition the maximum absorption capacity is 0.15 molCO2/molKOH. With a higher water content, the absorption capacity does not seem to be greatly affected. On the contrary, the absorption capacity decreases if the water content diminishes, becoming negligible when the system loses more than 15% of water with respect to its total weight.
Anyway, from the analysis of the system, it does not seem that the water has a special role in determining the structure of the mixture. So, there is no reason to believe that the decrease in absorption capacity is due to some change in the chemical structure of the system. More likely, it is the effect that water has to the viscosity of a mixture that reduces the
59
absorption capacity. Indeed, the latter does not seem to grow linearly with the viscosity, meaning that over a certain fluidity of the system the absorption remains the same. A higher viscosity, on the contrary, impedes the contact between the mixture and the atmosphere, reducing the absorption rate. Indeed, if the system was not regenerated by addition of water to the original weight, the CO2 absorption capacity drastically decreased until none was absorbed at all. This can be a problem, since at every cycle of desorption part of the water will evaporate together with the desorbed CO2. On the other hand, the system proved to be very stable during the absorption/desorption cycles, showing complete regeneration independently of how intense the desorption process might be or how many cycles are performed.
Anyway, the TGA of this system (Figure 29) could suggest an interesting solution to this inconvenient. This analysis shows a first major weight loss until about 100°C and this is believed to be caused primarily by water evaporation together with the release of the absorbed CO2. The weight loss at higher temperature is instead due to ethylene glycol evaporation and thermal degradation of the boron species. The analysis is conducted under a constant N2 flow that partially replaces the function of the vacuum applied during desorption, reducing the partial pressure of water, CO2 and EG. From this observation it can be guessed that a stream of water vapour could, in principle, work as a source of heat and hydration, facilitating the release as the N2 flow seemed to do. In so doing, desorption could be achieved simultaneously with regeneration of the system. In this way, it would be necessary a further step of separation of the water from the CO2 which it does not require particular energy (as a matter of fact, it is a step present in the
60
current plants for carbon capture) and the water could be recycled for future regeneration cycles. This hypothesis has not been investigated due to lack of time, but it is currently a proposal based on the data obtained. Besides, in the literature it is present a process of desorption based on the addition of water called “moisture swing adsorption” which, although it concerns different systems from the ones herein described, could have some interesting similarities.61
Finally, the melting points of the mixtures have been measured to further elucidate their nature and tentatively classify them as DESs or not. By this investigation, it is obvious that a major drop in the melting point occurs in all the systems making them eutectic mixtures but if this drop is due to the formation of a hydrogen bond network is not clear. From the characterization of these systems, it seems that the water content is maybe too high to consider them as DESs. Nevertheless, these systems are still liquids when are significantly dehydrated and, moreover, the KOH/EG is a viscous liquid even at the molar ratio 1:1. So, it is not to be excluded the hypothesis of being DESs but, if they are, with the water content of the initial composition are more likely to be considered as hybrid solution.62 Further analysis, in particular on the internal structure of these mixture, should be conducted to answer this question.
61
Electrochemical cycle
Scheme 9 shows the initially designed cycle of absorption/desorption for the system KOH/EG.
Scheme 9:
𝐾++ 𝑂𝐻−+ 𝐶𝑂2 → 𝐾+ + 𝐻𝐶𝑂3−
2𝐻𝐶𝑂3− + 𝐻2𝐶2𝑂4 → 𝐶2𝑂42−+ 2𝐶𝑂2
𝐶2𝑂42−+ 𝐻𝑂2− + 𝐻2𝑂 → 2𝐶𝑂2+ 3𝑂𝐻−
As already discussed in the introduction the latter reaction was not favoured and an electrochemical way to oxidize the oxalate had to be investigated to substitute it; the process is outlined in Scheme 10::
Scheme 10:
𝐶2𝑂42− → 2𝐶𝑂2+ 2𝑒−
2𝐻2𝑂 + 2𝑒− → 𝐻2 + 2𝑂𝐻−
𝐶2𝑂42−+ 𝐻2𝑂 → 2𝐶𝑂2 + 𝐻2+ 2𝑂𝐻−
As shown in the results, the feasibility of this process has been tested, at least in aqueous solution, with the application of a voltage around 1.00 V. In these preliminary tests, the ethylene glycol has been replaced with deionized water because, even though it proved to increase the CO2 absorption (a solution of KOH/H2O 1:3 has shown a maximum absorption of 0.48 molCO2/molKOH), most probably by the formation of the carbonated glycol, its influence in the electrochemical process is yet to be investigated. Furthermore, the proposed cycle could be adopted also without replacing the water with EG, as the heat capacity of the former would not be an issue anymore. Anyway, the electrical
