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Steel et al 2013_Conversion of CO2 into mineral carbonates


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Steel et al 2013_Conversion of CO2 into mineral carbonates

  1. 1. Conversion of CO2 into mineral carbonates using a regenerable buffer to control solution pH Karen M. Steel ⇑ , Kimia Alizadehhesari, Reydick D. Balucan, Bruno Bašic´ School of Chemical Engineering, University of Queensland, St. Lucia, Queensland 4072, Australia h i g h l i g h t s Use of a regenerable buffer (tertiary amine) is studied to enable mineral carbonation. Tertiary amines complex protons to give a pH of 8.2 which enables MgCO3 precipitation. Increasing temperature from 18 to 85 °C decreases pH by 2.5 pH units. Higher temperatures 85 °C might enable low pHs needed for Mg–silicate dissolution. a r t i c l e i n f o Article history: Received 25 October 2012 Received in revised form 14 March 2013 Accepted 16 April 2013 Available online 30 April 2013 Keywords: Mineral carbonation CO2 sequestration CO2 mineralisation Regenerable buffer a b s t r a c t The barrier that is currently stalling the rapid conversion of magnesium silicate deposits into magnesium carbonate as method for storing CO2 is considered to be the difference in pH needed for magnesium dis- solution from the silicate and magnesium precipitation as the carbonate, whereby rapid dissolution requires a low pH of around 1 while rapid precipitation requires a considerably higher pH of around 8. This paper investigates a novel concept which is to use a tertiary amine to bind with protons and raise the pH to around 8 and to then regenerate the amine through the use of heat due to the strength of the amine–proton bond decreasing with increasing temperature. This approach provides the low pH and high temperature that is needed for Mg dissolution and the high pH need for carbonate precipitation. The amine can be thought of as a regenerable buffer. Dissolution of Mg from serpentine has been found to be favourable with a solids to solution volume of more than 50 g/L to enable a low pH, and with temperatures close to the boiling point of the solution. The pH needed for magnesium carbonate precipitation was found to be approximately 8.2. Both triethyl- amine and tripropylamine were found to be capable of achieving this at 18 °C. Yields of around 20– 40 wt.% carbonate were achieved using residence times of approximately 1 h. The pH swing for the ter- tiary amines was found to be approximately 2.5 pH units between 5 and 85 °C, suggesting that an amine capable of achieving a pH of 8.2 at low temperature generates a pH of 5.7 in solution when heated to 85 °C. Further work will examine whether the lower pH values needed for serpentine dissolution can be achieved by heating the protonated amine to higher temperatures. Ó 2013 Elsevier Ltd. All rights reserved. 1. Introduction This paper investigates a novel concept to enable the conversion of CO2 into mineral carbonates using an aqueous route. The barrier that is currently stalling conversion is pH control. The first part of this introduction outlines why pH is critical, the second part out- lines what work has been done in the field of converting CO2 into mineral carbonates using an aqueous route, and the third part introduces the novel concept behind this study. 1.1. Thermodynamic modelling of the carbonate system In order for a mineral carbonate, such as MgCO3, to form the concentrations of CO2À 3 and Mg2+ in solution must be high enough for Eq. (1) to be satisfied [1]. ½Mg2þ Š½CO2À 3 Š 3:46  10À8 ð1Þ The solubility of CO2 in aqueous solution and the dissociation of carbonic acid to form bicarbonate and carbonate anions can be de- scribed by Eqs. (2)–(5) [1–4]. PPCO2 ðatmÞ ¼ 29:36½H2CO3Š ð2Þ ½HCOÀ 3 Š ¼ ½H2CO3Š=ð2:249  106 ½Hþ ŠÞ ð3Þ 0016-2361/$ - see front matter Ó 2013 Elsevier Ltd. All rights reserved. ⇑ Corresponding author. Tel.: +61 733653977; fax: +61 733654199. E-mail address: (K.M. Steel). Fuel 111 (2013) 40–47 Contents lists available at SciVerse ScienceDirect Fuel journal homepage:
  2. 2. ½CO2À 3 Š ¼ ½HCOÀ 3 Š=ð2:133  1010 ½Hþ ŠÞ ð4Þ ½Hþ Š ¼ ½HCOÀ 3 Š þ 2½CO2À 3 Š ð5Þ where 29.36 is the Henry’s law constant (dm3 atm molÀ1 ) for CO2 in water at 25 °C [2]. Here, [H2CO3] includes both dissolved molecular CO2 and molecular H2CO3. Eqs. (3) and (4) shows that CO2À 3 ions only start to become significant in solution when the pH is above about 8 (see Fig. 1). It could be thought that if the partial pressure of CO2 was raised the concentration of all species in solution (H2CO3, HCOÀ 3 and CO2À 3 ) would increase, and therefore a raised con- centration of CO2À 3 would result. However, solving Eqs. (2)–(5) for various partial pressures of CO2 (PPCO2 ) shows that the concentra- tion of CO2À 3 in solution stays very low at around 5  10À11 M de- spite the partial pressure of CO2 rising to 50 atm (see Fig. 2). This is due to the H+ ions, forming from the dissociation of H2CO3, always suppressing the formation of CO2À 3 (see Eqs. (3) and (4)). Only the concentration of HCOÀ 3 becomes appreciable at high CO2 partial pressures. Given the solubility limit of MgCl2 in solution (55 g/100 ml), the maximum possible concentration of Mg2+ is approximately 6 M Mg2+ and the corresponding carbonate concentration required to precipitate MgCO3 is 5  10À9 M. Given that 6 M Mg2+ is difficult to achieve as it is close to the solubility limit, a more realistic con- centration of 0.6 M Mg2+ would require a CO2À 3 concentration of 5  10À8 M CO2À 3 . Assuming a partial pressure of CO2 of 1 atm, Eqs. (2)–(4) can be solved to establish a relationship between car- bonate concentration and pH. Fig. 3 shows this relationship and shows that when the carbonate concentration is 5  10À8 M the equilibrium pH is approximately 5.5. This means that in order to precipitate MgCO3 the pH needs to exceed 5.5. 1.2. Work to date on converting CO2 into mineral carbonates using the aqueous route Bond and co-workers [5–7] looked at the potential of using an enzyme, called carbonic anhydrase, to convert power station CO2 into CaCO3. Coral reefs use carbonic anhydrase to assist in the pro- duction of CO2À 3 . The enzyme was found to catalyse the formation of CO2À 3 , however, although the catalyst enables the rapid forma- tion of CO2À 3 , there is also a simultaneous rapid drop in solution pH to approximately 4, which prevents the precipitation of carbon- ates. In the work of Bond et al. [5–7] and in recent work by Mirja- fari et al. [8], Ozdemir [9] and Rayalu and co-workers [10,11], precipitation of carbonates was only possible if a buffer was added to the solution in order to complex H+ ions and keep the pH high. The buffer used was tris(hydroxymethyl)aminomethane or ‘Tris’, which is commonly used to maintain pH at around 8. This buffer could not be continuously used in a large scale process, as it cannot be regenerated, which presents a major obstacle to the further development of this approach to sequestering CO2. Coral reefs have the ocean as a giant buffer for the H+ that they generate. It is worthwhile to note that Mirjafari et al. [8] found that calcium car- bonate did not precipitate when they used carbonic anhydrase with no buffer, but found precipitation to take place when they used the buffer with no carbonic anhydrase. Instead of converting salt solutions into carbonates, researchers have also looked at converting Mg–silicate minerals into carbon- ates in order to use the neutralising capacity of the mineral [12]. O’Connor and co-workers [13,14] found they could convert 34% of serpentine (Mg3Si2O5(OH)4) to MgCO3 using aqueous CO2 at 115 atm, 185 °C and a residence time of 24 h. The pH generated in solution would be around 3. The high temperature would be assisting the kinetics of Mg dissolution. Presumably the pH in- creases as dissolution proceeds and this aids precipitation of car- bonate. However, as the pH increases the dissolution rate of Mg would drop to negligible levels. The silica left behind and carbon- ate forming would also hinder further Mg dissolution/carbonation. O’Connor et al. found that artificially adding NaHCO3 to shift the equilibria in favour of a higher CO2À 3 to H+ ratio aided carbonation. pH 1 2 3 4 5 6 7 8 9 10 11 12 13 14 molfraction 0.0 0.2 0.4 0.6 0.8 1.0 H2 CO3 HCO3 - CO3 2- Fig. 1. Equilibrium distribution of H2CO3, HCOÀ 3 and CO2À 3 species in solution. CO2 partial pressure (atm) 10 -11 10 -10 10 -9 10 -8 10 -7 10 -6 10 -5 10 -4 10 -3 10 -2 10 -1 10 0 10 1 10 2 Concentration(M) 10 -12 10 -11 10 -10 10 -9 10 -8 10 -7 10 -6 10 -5 10 -4 10 -3 10 -2 10 -1 10 0 10 1 H2 CO3 HCO3 - CO3 2- Fig. 2. Equilibrium concentration of H2CO3, HCOÀ 3 and CO2À 3 in water as a function of CO2 partial pressure (derived from Eqs. (2)–(5)). pH 1 2 3 4 5 6 7 8 9 10 11 CO3 2- concentration(M) 10 -15 10 -14 10 -13 10 -12 10 -11 10 -10 10 -9 10 -8 10 -7 10 -6 10 -5 10 -4 10 -3 10 -2 10 -1 10 0 10 1 0.1atm 1 atm 10 atm 100 atm Fig. 3. Equilibrium concentration of CO2À 3 as a function of pH for CO2 partial pressures of 0.1, 1, 10 and 100 atm. K.M. Steel et al. / Fuel 111 (2013) 40–47 41
  3. 3. It follows that the best way to enable carbonation might be to have a two stage process where the first stage is optimised for Mg dissolution (high temperature and low pH) and the second is optimised for Mg carbonation (high pH). This need has been recog- nised as pH swing [15]. pH swing requires the removal of H+ ions from solution, which is usually achieved with soluble oxides/ hydroxides, however, these come from the calcination of carbon- ates which is obviously not possible. 1.3. Novel concept for converting CO2 into mineral carbonates using the aqueous route This paper investigates the use of weakly basic tertiary amines for complexing H+ ions generated when CO2 is added to solution and therefore enabling carbonate precipitation from a variety of salt solutions. The acid dissociation constant for protonated ter- tiary amines varies with temperature and so the approach is to use ‘‘temperature swing’’ to regenerate the amine, whereby the loaded amine is heated to strip the acid off. This concept has re- cently been patented [16]. Conventional CO2 capture technologies involve absorbing CO2 into a mixture of primary and tertiary amines, including monoeth- anolamine (MEA) and methyldiethanolamine (MDEA) respectively. The reason for this is that while the primary amine forms a strong carbamate bond and therefore enables a high CO2 loading due to the strength of the bond considerable energy is needed to break it and regenerate the MEA (approximately 3–5 MJ/kg CO2 [17]). In order to reduce the energy load tertiary amines are blended. Ter- tiary amines do not bond with CO2 but rather strip the solution of H+ ions thus driving the formation of HCOÀ 3 and CO2À 3 ions in solu- tion. The loading of CO2 in solution as these ions is much lower but the energy needed to regenerate the MDEA is much less, such that the total energy needed to regenerate the MEA/MDEA mixture is around 1–3 MJ/kg CO2 [17]. The introduction of MDEA means that taller towers are needed for a given separation efficiency and so a balance between capital and operating (energy) cost must be struck. The idea put forward here is to use tertiary amines alone to con- vert CO2 into mineral carbonates. It is thought that the loading of CO2 can be higher than that for conventional CO2 capture using MDEA because the CO2À 3 ions that form leave the solution as solid carbonate, thus providing a stronger driving force for the capture of CO2 into solution. The advantage of this concept for CO2 sequestration is twofold. Firstly, the energy needed for the process might be considerably lower than that needed for conventional CO2 capture because less energy is needed to regenerate the amine and compression of the CO2 (needed for storage) is not necessary. Secondly, the CO2 would be locked up in a mineral form that is known to be stable for mil- lions of years, which is an important consideration given the scale of CO2 that needs to be stored. Fig. 4 shows a simplified diagram of the novel concept that is being explored. In stage 1, acid loaded amine is heated to $100 °C and contacted with a Mg silicate such as serpentine ((Mg, Fe)3Si2O5(OH)4). At the high temperature the acid (HCl) dis- sociates from the amine, thereby providing a low pH capable of dissolving Mg out of the serpentine. The Mg depleted serpentine is separated by density and/or filtration and the solution, contain- ing MgCl2 and regenerated amine, passes to stage 2. In stage 2, the solution is cooled and the flue gas containing CO2 is sparged through. At the low temperature acid generated by the dissociation of H2CO3 plus excess acid generated in stage 1 are complexed by the tertiary amine which causes the pH to rise and consequently the CO2À 3 concentration to rise to a level that is sufficient to begin interacting with Mg2+ and precipitating MgCO3. The MgCO3 is sep- arated by density and/or filtration and the solution, containing acid loaded amine is recycled to stage 1. The overall reaction taking place in stage 1 is as follows: Mg3Si2O5ðOHÞ4 þ 6R3NHCl ! 3MgCl2 þ 6R3N þ 5H2O þ 2SiO2 ðR1Þ And the overall reaction taking place in stage 2 is as follows: MgCl2 þ CO2 þ H2O þ 2R3N ¡ MgCO3 þ 2R3NHCl ðR2Þ The concept shown in Fig. 4 can be used in a variety of different ways by a variety of different industries, and is not locked into one industrial application. The concept could be used for the treatment of magnesium sil- icate deposits as described above. It is worthwhile to note that many magnesium silicate deposits contain significant levels of valuable heavy metals such as Ni, called Ni laterites, and the pro- cess could therefore have the dual operation of Ni extraction com- bined with CO2 sequestration. It is known that treating Ni laterites with acid to dissolve the Mg enables a greater amount of Ni to be extracted into solution. Secondly, the concept could be used for salt mining operations. Chloride and sulphate salt deposits are mined for KCl or K2SO4 to be used as fertiliser. By using solution mining the salt comes up hot and is cooled to separate the potassium salts. The refuse salt represents a waste that is generally deposited back into the re- serve. The salt solution could be processed using the sequestration technology shown in Fig. 4 to form a mixture of carbonates and bicarbonates that are deposited back into the reserve for longterm CO2 storage. The heat that must be taken out of the solution as it comes to the surface could be used to provide the heat needed for stage 1. Using the concept for salt solutions means that a by- product of the process is acid, either hydrochloric or sulphuric acid. Therefore, the scale of the operation would need to be matched with HCl or H2SO4 needs in the oil, chemical and mineral sectors. This paper presents our work to date on this novel concept. Fig. 4. Simplified block flow diagram of the proposed CO2 sequestration technology. 42 K.M. Steel et al. / Fuel 111 (2013) 40–47
  4. 4. 2. Experimental 2.1. Serpentine sample The serpentinite sample (Mg3Si2O5(OH)4) used for this study was obtained from a naturally occurring deposit in northern Queensland, Australia. It was initially ground by hand in a pestle and mortar and then in a laboratory attrition mill. The ground ser- pentinite was sieved with ASTM standard sieves to obtain particles with a diameter of 57 lm. Australian Laboratory Services (ALSs) performed the elemental analysis via alkali fusion, acid digestion and inductively coupled plasma-atomic emission spectroscopy (ICP-AES) of the resulting solution. The loss on ignition at 1000 °C (LOI1000) was also performed using a TGA furnace. The re- sults of ALS’s analysis based on their method ME-ICP85 (Silicates by Fusion, ICP-AES) and ME-GRA05 (H2O/LOI by TGA Furnace) are summarised in Table 1. Mineral composition was first probed via X-ray diffraction anal- ysis using a PANanalytical XPERT-PRO diffractometer with Cu Ka target (k = 0.15406 nm) at room temperature. Measurements were made in a step scan mode (0.1°/step) over the 2h range of 10–90°. Phase matching of the X-ray powder diffraction (XRPD) pattern of the serpentinite sample against the International Centre for Dif- fraction Data (ICDD) database suggested antigorite to be the pri- mary serpentine phase present. A calibrated TA Instrument SDTQ600 Thermogravimetric analyser-differential scanning calo- rimeter provided further mineral characterisation via thermogravi- metry–derivative thermogravimetric analysis (TGA–DTG). Replicate runs were obtained for 10 mg of the À53 lm samples using alumina crucibles and heated from 30 °C to 1000 °C at a heat- ing rate (b) of 10 °C minÀ1 . A 10-min isothermal stage was em- ployed at 110 °C to determine the moisture content of the material and was found to contain 1.0 wt.%. The total mass loss of the dry sample (Dm105–850 °C) was determined as 11.5 ± 0.01 wt.%, which is in fair agreement with the analysis made by ALS (11.7 wt.%). Fig. 5 shows the TGA–DTG profile of the sample, where the characteristic serpentine doublet comprising the DTG temperature shoulder, Tsh, showing at 597 °C and the peak temperature, Tp, at 718 °C. Thermal analysis suggests that this particular serpentinite sample is fully serpentinized and contains antigorite as well as lizardite (antigorite + lizardite). The antigorite component displays its DTG peak temperature, Tp1ATG at 718 °C and its diagnostic peak, Tp2ATG at 747 °C. The shoulder at 701 °C is thought to indicate the lizardite component, Tp1LIZ. Based on XRPD and TGA-DTG analysis, we then refer to this sample as ‘‘serpentinite’’, rather than antigor- ite as this specimen also contains lizardite. 2.2. Magnesium dissolution Magnesium dissolution experiments were carried out using AR grade HCl (37 wt.%) and Millipore water. 0.5 g of sample and acid solution was mixed in a 250 ml spherical flat-bottom flask mounted on a magnetic stirrer/hotplate and equipped with a cold water condenser. The effects of temperature, concentration of HCl, residence time and acid solution volume were investigated. At the duration of the experiment, the residue was vacuum filtered, dried overnight and weighed. The pH of the filtrate solutions was mea- sured using a pH electrode. The solid remaining was analysed by ALS using the procedure described above to determine the extent of Mg dissolution. 2.3. Carbonate precipitation Carbonate precipitation experiments were carried out using a 250 ml Erlenmeyer flask open to the atmosphere. Food grade CO2 was injected into the solution through a sparger containing five holes of 2 mm diameter each. Pressure was regulated at 1.4 bar and flow was set at approximately 1 L/min using a rotameter. While sparging the solution with CO2, tertiary amine was added dropwise via a burette while simultaneously measuring pH via a pH electrode. The amines investigated were simple straight chain trialkylamines with increasing chain length, i.e. triethylamine, tri- propylamine, tributylamine, tripentylamine, trihexylamine. The solution was observed for the onset of precipitation. If a precipitate form, it was filtered through Whatman No. 1 filter paper, dried and weighed. Precipitated solids were analysed by XRD and ICP-AES for compound and elemental determinations, as described above. 2.4. Amine regeneration The ability of the amines to be regenerated and liberate bound acid was investigated via a series of titrations at various tempera- tures whereby the amines were added to a standardised solution of 0.1 M HCl (10 ml). The HCl was placed in a 3-neck round bottom flask which was immersed in a water bath to control temperature. A condenser was fitted vertically to the middle neck and a ther- mometer and pH electrode were inserted and sealed through each of the side necks. Amine was added through the opening at the top of the condenser using an automatic pipette. For experiments per- formed at 5 °C ice was added to the water bath. The tertiary amines studied were triethylamine, tripropylamine, tributylamine and tripentylamine. 3. Results and discussion 3.1. Magnesium dissolution The effect of HCl concentration on the dissolution of Mg from serpentine is shown in Fig. 6. All percentages are weight percent- ages. The residence time used was 3 h and the temperature was the boiling temperature of the solution ($100 °C). The stoichiome- tric amount of acid needed to dissolve all of the Mg according to R1 is approximately 0.12 M which gives approximately 40% extrac- tion, while a plateau of approximately 65% extraction is reached at around 0.5 M HCl. Fe and Al were also found to dissolve with similar extraction levels to those of Mg, showing that the elements do not appear to dissolve selectively with respect to HCl concentration. Because it is desirable to not have excess acid in solution to minimise the energy needed for amine regeneration a compromise between Mg extraction and solution pH must be struck. Fig. 6 shows the pH change as a function of HCl concentration. With twice the stoichiometric amount needed, the pH of the spent solu- tion is still $1. The effect of residence time on dissolution of Mg in 0.25 M HCl is shown in Fig. 7. Equilibrium is reached after approximately 3 h. The effect of temperature on the dissolution of Mg in 0.25 M HCl is shown in Fig. 8. For temperatures less than 50 °C only a small amount of Mg dissolves. At temperatures higher than 50 °C a linear Table 1 Chemical composition of the serpentinite sample. wt.%, ±0.01 MgO* SiO2 * Fe2O3 * Al2O3 * CaO* Ni* MnO* K2O* LOI1000 ** 39.3 44.1 7.38 1.03 0.35 0.24 0.10 0.04 11.7 * Values obtained via inductively coupled plasma-atomic emission spectroscopy on fused samples after acid digestion. Based on ALS’s ME-ICP85. ** Value obtained by heating the moisture free sample to 1000 °C using TGA fur- nace. Based on ALS’s ME-GRA05. K.M. Steel et al. / Fuel 111 (2013) 40–47 43
  5. 5. increase in Mg dissolution with respect to temperature is obtained, reaching approximately 65% at $100 °C and suggesting that at temperatures above 100 °C higher extraction efficiencies might be achieved. Extrapolation suggests that close to 100% extraction might be achieved at 140 °C. Experiments using a pressurised ves- sel to enable higher temperatures above 100 °C are recommended to confirm the extrapolated trend shown. As presented in the introduction, it is desirable to have a high Mg2+ concentration in solution as this reduces the concentration of CO2À 3 needed for MgCO3 precipitation. Experiments were per- Fig. 5. The TGA–DTG profile of the serpentinite sample used in this study with the characteristic peaks indicated. Fig. 6. Effect of HCl concentration on the dissolution of Mg and final pH (residence time 3 h, $100 °C, 0.5 g, 100 ml). Fig. 7. Effect of residence time on the dissolution of Mg (0.25 M HCl, solution and $100 °C, 0.5 g, 100 ml). Fig. 8. Effect of temperature on the dissolution of Mg (0.25 M HCl, residence time 3 h, 0.5 g, 100 ml). Fig. 9. Effect of solution volume on the dissolution of Mg (0.025 mols HCl, residence time 3 h, 0.5 g, $100 °C). 44 K.M. Steel et al. / Fuel 111 (2013) 40–47
  6. 6. formed keeping the amount of HCl the same and decreasing the solution volume from 100 ml. Twice the stoichiometric amount needed (0.024 mols) was chosen for the amount of HCl. Fig. 9 shows the effect of reducing the solution volume down to 10 ml. The extraction increases as the solution volume decreases reaching approximately 85% with only 10 ml of solution. The concentration of Mg in solution is approximately 0.35 M. This work has shown the importance of both acid concentration and temperature on the dissolution of Mg. It is recommended to operate with a solids to solution volume of more than 50 g/L for the extraction stage, a temperature close to the boiling tempera- ture of the solution or higher if using a pressurised vessel, a resi- dence time of 3 h and concentration no more than twice the stoichiometric amount needed for reaction. These conditions en- able high extractions of Mg approaching 100%. 3.2. Carbonate precipitation The extract solution from the experiment with 10 ml solution volume shown in Fig. 9 was used for carbonation. Tripropylamine was added dropwise. As the pH increased to approximately 5 a light brown precipitate formed which was found from elemental analysis to contain around 20.1 wt.% Fe, 10.8 wt.% Si and 5.5 wt.% Al and only 0.2 wt.% Mg. This product comes from the hydrolysis of Fe, Si and Al which dissolved during serpentine dissolution. The leaching studies had shown that the elements dissolved simul- taneously with Mg. After removing this precipitate further addi- tions of TPA raised the pH to approximately 8 at which point a white precipitate formed. This precipitate began forming within a few minutes. After bubbling CO2 for 45 min, the precipitate was recovered by filtration, dried and analysed. The weight was 0.20 g. XRD analysis indicated the formation of nesquehonite (MgCO3Á3H2O). Elemental analysis showed that the purity was high with a composition of 18.8 wt.% Mg, 0.2 wt.% Ca, 0.06 wt.% Fe and 0.01 wt.% each of Al and Si. The yield was 29.1 wt.% (i.e. 29.1 wt.% of the Mg extracted into solution was converted to the MgCO3 precipitate). The remaining Mg would be recycled to the first stage of the process. In order to study MgCO3 more precisely and the pH level needed for carbonate precipitation without the hindrance of other dis- solved elements model compound work was performed using Mg(OH)2. 0.126 g of Mg(OH)2 was dissolved in 50 ml of 0.0965 M HCl such that the acid was slightly in excess and gave a final pH of 2.08 after the Mg(OH)2 had completely dissolved. To this solu- tion 5 Â 10À4 mols of TPA was added, which is the amount needed to neutralise the excess acid. The pH rose to 9.59. CO2 was then bubbled through the solution and after 5 min the pH had decreased and stabilised at 4.62 and no precipitate had formed. A further addition of TPA was made (0.006 mols) and the pH stabilised at approximately 8.27. This addition of TPA is 1.5 times that needed for reaction 2. Over the next half an hour, CO2 was continually bub- bled through the solution. It was found that TPA need to be contin- ually added dropwise in order to maintain the pH at a level above 8. The total amount of TPA added (neglecting the initial 5 Â 10À4 - mols) was 0.021 mols and the final pH was 8.43. The solution was filtered and the solid recovered and air dried. The weight of the so- lid was 0.051 g and as found above, analysis indicated nesqueho- nite (MgCO3Á3H2O) to be the primary phase present. The yield was approximately 17% and the amine used was 5 times in excess. The above experiment was repeated with slower additions of TPA. A total amount of 0.0085 mols (twice excess) was added over a period of 1 h. The mass of solid recovered was 0.073 g (dried) which is a yield of 24 wt.%. The above experiment was repeated with triethylamine (TEA). 0.164 g of Mg(OH)2 was dissolved in 50 ml of 0.112 M HCl, which is the stoichiometric amount needed for complete dissolution. CO2 was bubbled which decreased the pH to 5.38. 0.0072 mols of TEA was added and the pH increased to 10.16 and simultaneously the solution became milky with precipitation. As CO2 addition con- tinued the pH decreased to 7.27 even though the amount of TEA added was 30% above that needed for reaction 2. A further addition of 0.0036 mols of TEA increased the pH to 9.32 initially but then it stabilised at 8.23. The solution was filtered to recover the solid, which had a mass of 0.133 g and therefore a yield of 34 wt.%. Tests with both tributylamine and tripentylamine did not yield precipitates, which is thought to be due to the pH generated by the amines not being high enough. It is possible that decreasing the temperature would enable precipitation to take place with these amines. These experiments have shown that the pH needed for magne- sium carbonate precipitation is approximately 8.2 and that trieth- ylamine and tripropylamine are capable of achieving this. An excess of amine has been found to be necessary to maintain the pH of 8 while CO2 is bubbled through the solution. So far yields of around 20–40 wt.% have been achieved. The reason why precip- itation did not occur at lower pH levels, such as the pH level of 5.5 predicted from the theoretical modelling work reported in the introduction, is thought to be due to the kinetics being too slow be- low 8.2. It was found with TEA that precipitation was within sec- onds when the pH was 10. Further experiments will investigate the kinetics of carbonate precipitation by sampling periodically, particularly during the early stages. 3.3. Amine regeneration Based on the serpentine dissolution work, to achieve dissolu- tion of Mg under reasonable conditions of a residence time of less than an hour and reaction time of 100–150 °C, the pH of the solu- tion needs to be approximately less than 1. Based on the carbon- ation work, the pH of the Mg rich solution needs to be raised to approximately 8.2. To examine the ability of tertiary amines to reversibly enable this change a series of titrations were performed. The dissociation constant for triethylamine at various tempera- tures has been published by Hamborg and Versteeg [18], whereby the pKa is 10.89 at 18 °C and decreases to about 9.17 at 90 °C. Fig. 10 shows these constants converted into a titration curve whereby the amine is being added to 0.1 M HCl. Our own titration points are also shown for comparison. The titration curve is ex- pressed this way as it mimics the real system. Total amine essen- Fig. 10. Titration curve for triethylamine against 0.1 M HCl at 18 °C and 90 °C generated from pKa data obtained from literature and at 18 °C generated from experiment. K.M. Steel et al. / Fuel 111 (2013) 40–47 45
  7. 7. tially means the amount of amine added expressed as a concentra- tion. At 20 °C the pH rises to approximately 11 while at 90 °C it rises to approximately 9. While a pH of 11 would assist with the precipitation of carbonates the pH of 9 at the higher temperature will not assist with the dissolution of Mg from serpentine. Figs. 11–13 shows our own titration curves for tripropylamine, tributylamine and tripentylamine at various temperatures. For tri- propylamine, the final pH is approximately 9.5 at 18 °C and 7.1 at 85 °C. For tributylamine, the final pH is approximately 8.6 at 5 °C and 6.0 at 85 °C. For tripentylamine, the final pH is approximately 6.5 at 5 °C and 4.0 at 85 °C. It follows that with a temperature rise from 5 to 85 °C, the change in final pH is approximately 2.5 pH units. This work shows that at elevated temperatures an acid loaded alkylamine with a long chain length behaves similarly to a weak acid and therefore might have the ability to dissolve Mg from serpentine. The equivalence point (point of highest gradient) was found for each titration curve from which pKa values were estimated. Table 2 shows results from this analysis and Fig. 14 shows the constants for tributylamine and tripentylamine as a function of temperature. Lines of best fit have been drawn through the data and extended to 135 °C. Fig. 14 also shows a line corresponding to the pKa values reported by Hamborg and Versteeg [18] for triethylamine. The lines obtained from this work appear to decrease more steeply than those for triethylamine, which may be due to experimental error associated with vapour losses from the system which could concentrate the protons and give lower pH values. Further experiments to obtain more accurate titration data including data at higher temperatures and pressures (100–150 °C and 1–5 bar) is planned for the future. The experiments at higher temperature will also involve treating serpentine with the regener- ated amine and acid mix to study the dissolution behaviour. These experiments are akin to treating serpentine with weak organic acids. There is currently a lack of studies on the behaviour of ser- pentine with weak organic acids particularly at high temperatures where the kinetics of dissolution is favourable. Teir et al. [19] have reported the behaviour of carboxylic acids alongside stronger acids however the studies were confined to 20 °C. Unlike strong acids which provide a high initial concentration of protons which de- creases as the mineral dissolves, weak acids provide a low concen- tration that is maintained as the mineral dissolves, and protons are consumed in the acid-base reaction thereby driving the dissocia- tion of more protons from the weak acid. Fig. 11. Titration curve for tripropylamine against 0.1 M HCl at 18 and 85 °C. Fig. 12. Titration curve for tributylamine against 0.1 M HCl at 5, 18 and 85 °C. Fig. 13. Titration curve for tripentylamine against 0.1 M HCl at 5, 18 and 85 °C. Table 2 Calculated pKa values for tertiary amines at 5, 18 and 85 °C. Amine pKa 5 °C 18 °C 85 °C Tripropylamine nd 9.47 7.32 Tributylamine 9.84 8.44 6.32 Tripentylamine 6.81 5.80 3.98 nd: Not determined. Fig. 14. Expected trend for pKa as a function of temperature for tributylamine and tripentylamine with lines of best fit, and pKa line for triethylamine as derived from Hamborg and Versteeg [18]. 46 K.M. Steel et al. / Fuel 111 (2013) 40–47
  8. 8. If the regenerated amines enable high degrees of serpentine dis- solution, experiments will move to examining the extent that the amines can be continually recycled. 4. Conclusion The best conditions for the dissolution of Mg from serpentine have been found to be a solids to solution volume of more than 50 g/L to enable a high proton concentration. The amount of acid should be no more than twice the stoichiometric amount needed for reaction. Reaction temperature should be as high as possible, close to the boiling temperature of the solution or higher (100– 150 °C) if using a pressurised vessel. These conditions combined with a residence time of 3 h are able to dissolve approximately 85% of the Mg in serpentine. These experiments have shown that the pH needed for magne- sium carbonate precipitation is approximately 8.2 and that trieth- ylamine and tripropylamine are capable of enabling this at 18 °C. It appears that an excess of amine is needed to maintain the pH of 8 while CO2 is bubbled through the solution. So far yields of around 20–40 wt.% have been achieved for tripropylamine using residence times of approximately 1 h. Precipitation occurred more rapidly for triethylamine owing to the higher pH generated. The association of tertiary amines with HCl has been found to decrease with increasing temperature such that there is a differ- ence of approximately 2.5 pH units between 5 and 85 °C. This means that an amine capable of achieving a pH of 8.2 at low tem- perature generates a pH of 5.7 in solution when heated to 85 °C. While this is not low enough to provide a high rate of serpentine dissolution it is thought that increasing the temperature beyond 85 °C may yield pH levels capable of dissolving high levels of Mg, particularly given that high temperatures aid the kinetics of disso- lution. 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