CO2 Separation Using Regenerable Magnesium Solutions Dissolution

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CO2 SEPARATION USING REGENERABLE MAGNESIUM SOLUTIONS

DISSOLUTION, KINETICS AND VLSE STUDIES

A thesis submitted to the Graduate School of the University of Cincinnati in partial fulfillment of the

requirements for the degree of

Master of Science

In the Chemical Engineering Program of School of Energy, Environmental,

Biological and Medical Engineering

(College of Engineering and Applied Science)

Hari Krishna Bharadwaj

B.Tech, Sri Venkateswara College of Engineering (Anna University), 2009

Committee Members:

Dr. Joo-Youp Lee (Chair)

Dr. Soon-Jai Khang

Dr. Tim C. Keener ABSTRACT

Fossil fuel power plants are responsible for a considerable amount of total anthropogenic CO2 emissions throughout the world. CO2 capture and sequestration from point source emitters like fossil fuel power plants is essential to mitigate the harmful effects of greenhouse gases. A novel post-combustion CO2 capture technique has been proposed, where CO2 from flue gases can be absorbed by magnesium hydroxide (Mg(OH)2) slurry at 52 °C in a column followed by a regeneration step in a stripper. CO2 absorption in a Mg(OH)2 slurry solution involves physical absorption of CO2 gas into the aqueous phase and subsequent chemical reactions in the aqueous and solid phases. Three critical phenomena influencing the CO2 absorption process, namely: the dissolution of Mg(OH)2 for continuous CO2 absorption, formation of an unwanted by-product magnesium carbonate (MgCO3), and vapor-liquid-solid equilibrium (VLSE) of the system were studied. The dissolution of Mg(OH)2 and the release of magnesium ions into the solution to maintain a level of alkalinity is a crucial step in the CO2 absorption process. The dissolution process was modeled using the shrinking core model and surface chemical reaction was found to be the rate controlling mechanism. The formation of MgCO3 will reduce a regenerative capacity of Mg(OH)2 solvent. The impact of pH control and temperature on the kinetics of magnesium carbonate formation was also studied. The vapor-liquid equilibrium data for a solvent-CO2 system is essential for the design and operating conditions of an absorber and a desorber. It can be used to determine an amount of Mg(OH)2 feed for CO2 absorption and also to determine the operating conditions for CO2 separation and Mg(OH)2 regeneration from a CO2 desorber. These three individual studies will provide a fundamental understanding of CO2 absorption in a

Mg(OH)2 slurry solution and basic engineering data for the design and operation of a Mg(OH)2- based CO2 absorption system.

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ACKNOWLEDGEMENTS

I would like to express my sincere gratitude and thanks to the following individuals and groups, who have contributed immensely during the course of my journey to obtain a M.S. degree in

Chemical Engineering at the University of Cincinnati:

To Dr. Joo-Youp Lee, the most sincere and committed professor I have ever interacted with. I wish to emulate your passionate zeal towards work and thank you for your support, guidance and patience.

To Dr. Timothy Keener, for his useful suggestions during group meetings and agreeing to grace my committee.

To Dr. Soon-Jai Khang, for agreeing to grace my committee.

To Padmini Lingaraju, for all the patient advice and helping me out in the lab. To Jinsoo Kim,

Zhouyang Liu, Lei Cheng, Tongyan Li and all other lab members of both Dr. Lee’s and Dr.

Keener’s group, for their support and cooperation.

To my friends in Cincinnati, especially the Punyabhoomi group, for making this journey, a fun filled and joyous one. To Bhavana, for being there through all the crests and troughs of this journey.

Most importantly, to my parents, Dr. Bharadwaj and Mrs. Vasanthi Bharadwaj for constantly supporting and encouraging me in all aspects. To Gowri, for being the best sister in the world.

And finally, to the University of Cincinnati, US-EPA and US-DOE for supporting my education/research, for which I am forever indebted.

TABLE OF CONTENTS

ABSTRACT ...... ii

CHAPTER 1. LITERATURE SURVEY ...... 1

1.1 Need for CO2 Capture and Sequestration...... 1

1.2 Various Technologies Involved in CO2 Capture and Sequestration ...... 2

1.3 Challenges facing Post-Combustion CO2 Capture ...... 6

1.4 Post-Combustion CO2 Capture Pathways ...... 9

1.5 Post-Combustion CO2 Capture using Absorption ...... 9

1.5.1 Physical Absorption using Solvents ...... 9

1.5.1.1 Sulfolane ...... 10

1.5.1.2 Selexol ...... 11

1.5.1.3 Rectisol ...... 11

1.5.1.4 Fluor Process ...... 12

1.5.2 Chemical Absorption using Solvents ...... 12

1.5.2.1 Amine-Based Solvents ...... 13

1.5.2.2 Carbonate-Based Solvents ...... 14

1.5.2.3 Ammonia ...... 15

1.5.2.4 Dual Alkalis ...... 16

1.5.2.5 Ionic Liquids ...... 16

1.5.3 Absorption using Carbonation/Calcination Cycles ...... 17

1.6 Post-Combustion CO2 capture using Adsorption ...... 18

1.7 Post-Combustion CO2 Capture using Membranes ...... 19

1.7.1 Polymeric Membranes ...... 20

1.7.2 Zeolitic Membranes...... 21

1.7.3 Membranes in Conjunction with Chemical Solvents ...... 22

1.8 Post-Combustion CO2 Capture using Biological Methods ...... 23

1.9 Other Technologies for CO2 Capture ...... 24

1.9.1 Cryogenics ...... 24

1.9.2 Metal Organic Frameworks ...... 24

1.10 Conclusion ...... 25

CHAPTER 2. MAGNESIUM HYDROXIDE AS POST-COMBUSTION SOLVENT FOR CO2 ABSORPTION ...... 26

2.1 Magnesium Hydroxide for CO2 Capture...... 26

2.2 Process Description ...... 27

2.3 Advantages of Using Magnesium Hydroxide for CO2 Capture ...... 28

2.4 Magnesium Hydroxide Recovery from Magnesium-Enhanced Wet FGD ...... 29

2.5 Chemistry of the Absorption / Desorption Process...... 31

CHAPTER 3. DISSOLUTION RATE STUDIES OF MAGNESIUM HYDROXIDE ...... 33

3.1 Introduction ...... 33

3.2 Motivation in Estimating the Dissolution Rate of Mg(OH)2 ...... 36

3.3 Dissolution Rate Studies – Experimental Techniques ...... 37

3.4 Materials and Methods ...... 38

3.4.1 Experimental Setup ...... 38

3.4.2 Estimation of Solid Dissolution Flux ...... 39

3.5 Results and Discussion ...... 40

3.5.1 Effect of pH ...... 42

3.5.2 Effect of Temperature...... 43

3.5.3 Effect of Agitation Speed ...... 45

3.5.4 Experimental Model ...... 46

3.6 Kinetic Analysis and Shrinking Core Model ...... 47

3.6.1 Selection of Rate Controlling Mechanism ...... 47

3.6.2 Conversion-Time Equations and Determination of Apparent Rate Constant ...... 50

3.6.3 Order of Reaction and Overall Activation Energy for Mg(OH)2 Dissolution ...... 53

3.6.4 Modified Rate of Dissolution Expression Based on the Shrinking Core Model ...... 56

3.6.5 Effect of Decreasing Particle Size During the Dissolution Reaction...... 57

3.6 Conclusion ...... 62

CHAPTER 4. KINETICS OF MgCO3 FORMATION...... 63

4.1 Introduction ...... 63

4.2 Experimental Setup and Procedure ...... 65

4.2.1 Solid Analysis using the Thermo-gravimetric Analyzer - Mass Spectrometer...... 66

4.3 Results and Discussion ...... 68

4.3.1 No pH control and Temperature of 52 C ...... 69

4.3.2 pH control at 8.6 and temperature of 52 C ...... 71

4.3.3 No pH control and temperature of 65 C ...... 73

4.3.4 pH control at 8.6 and temperature of 65 C ...... 76

4.3.5 No pH control and room temperature [24 °C] ...... 78

4.3.7 pH control at 8.6 and room temperature [24 °C] ...... 80

CHAPTER 5. VAPOR-LIQUID-SOLID EQUILIBRIUM (VLSE) STUDIES ...... 83

5.2 Experimental Setup and Procedure ...... 86

5.2.1 Experimental Procedure ...... 87

5.2.2 Calculation Procedure ...... 88

5.3 Results and Discussion ...... 88

5.3.1 Room Temperature Data ...... 88

5.3.2 VLSE Data at 40 oC and 52 oC ...... 91

5.4 Conclusion ...... 95

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CHAPTER 6. CONCLUSION ...... 96

REFERENCES ...... 98

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LIST OF FIGURES

Figure 1. Block flow diagrams of the three technological pathways for CO2 capture.

Figure 2. Diagram of CO shift reaction in post combustion CO2 capture.

Figure 3. Oxyfuel technology for CO2 capture and storage.

Figure 4. Typical post combustion CO2 capture process

Figure 5. Additional energy requirements for CO2 capture.

Figure 6. Post Combustion capture pathways.

Figure 7. Process flow diagram of the proposed setup.

Figure 8. Schematic of magnesium hydroxide recovery process.

Figure 9. Different techniques for estimating dissolution rate.

Figure 10. Schematic for the dissolution rate experiment.

Figure 11. Sample pH vs. time (pH = 8.6, temperature = 52 oC and rpm = 700)

Figure 12. Weight of HCl solution vs. time (pH = 8.6, temperature = 52 oC and rpm = 700)

Figure 13. ln Ns (Solid Dissolution Flux) vs. pH.

Figure 14. Solid dissolution flux vs. temperature.

Figure 15. Arrhenius plot of ln Ns (solid dissolution flux) vs. 1/T.

Figure 16. ln Ns (solid dissolution flux) vs. ln ω.

Figure 17. Comparison between the experimental and predicted values of the dissolution flux.

Figure 18. Rate controlling mechanism at temperature = 52 oC / 700 rpm / HCl stroke rate = 15.

Figure 19. Rate controlling mechanism at temperature = 52 oC / 700 rpm / HCl stroke rate = 30.

1/3 o Figure 20. Sample plot of (1-(1-XB) ) vs. Time (min) for pH = 8.6 and temperature = 52 C.

o Figure 21. Activation energy calculation using kr at pH = 8.6 and temperature = 52 C.

Figure 22. Plot of ln kr vs ln Cal to determine n and k” at different temperatures.

Figure 23. Plot of ln k’’ vs. 1/T to determine the overall activation energy.

Figure 24. Comparison of measured data with SCM model at pH = 8.6 and temperature = 52 oC

Figure 25. Particle size distribution for sample taken at 0 min.

Figure 26. Average particle size vs. time using SCM model and experimental data.

Figure 27. Fraction of unreacted area vs. time.

Figure 28. Weight of HCl consumed vs. time for pH = 8.6, temperature = 52 οC for 60 min.

Figure 29. Plot showing the effect of particle size on the rate controlling mechanism.

Figure 30. Distribution of dissolved magnesium species as a function of pH.

Figure 31. Distribution of dissolved carbonate species as a function of pH.

Figure 32. TGA curve for MgCO3.

Figure 33. Experimental setup for kinetic analysis.

Figure 34. Calibration curve for CO2 measurement using the TGA-MS.

Figure 35. Sample TGA-MS curve for pH control and temperature 65 oC at 0 min.

Figure 36. Sample TGA-MS curve for pH control and temperature 65 oC at 15 min.

Figure 37. Carbon content vs. time for no pH control at 52 °C.

Figure 38. pH vs. time graph for no pH control at 52 °C.

Figure 39. Carbon content vs. time for pH control at 52 °C.

Figure 40. pH vs. time for pH control at 52 °C.

Figure 41. Carbon content vs. time for no pH control at 65 °C.

Figure 42. pH vs. time for no pH control at 65 °C.

Figure 43. Carbon content vs. time for pH control at 65 °C.

Figure 44. pH vs. time for pH control at 65 °C.

Figure 45. Carbon content vs. time for no pH control at 24°C.

Figure 46. pH vs. time for no pH control at 24 °C.

Figure 47. Carbon content vs. time for pH control at 24 °C.

Figure 48. pH vs. time for pH control at 24 °C.

Figure 49. Schematic diagram of the experimental setup.

Figure 50. VLSE data for Mg(OH)2-CO2-H2O at room temperatures.

Figure 51. VLSE data for Mg(OH)2-CO2-H2O at higher temperatures.

o Figure 52. VLSE data for Mg(OH)2-CO2-H2O at 52 C and different concentrations.

Figure 53. Comparison of CO2 loading capacities of various solvents.

LIST OF TABLES

Table 1. Energy & efficiency penalties for various CO2 capture techniques

Table 2. Summary of dissolution rate results

Table 3. Mass balance for no pH control and temperature of 52 oC

Table 4. Mass balance for pH control at 8.6 and temperature at 52 oC

Table 5. Mass balance for no pH control and temperature of 65 oC

Table 6. Mass balance for pH control at 8.6 and temperature of 65 oC

Table 7. Mass balance for no pH control and temperature of 24 oC

Table 8. Mass balance for pH control and temperature of 24 oC

Table 9. Literature review of VLE estimation

Table 10. CO2 equilibrium data for supernatant Mg(OH)2 solution

Table 11. CO2 equilibrium data for 0.1 M Mg(OH)2 solution

Table 12. CO2 equilibrium data for 1 M Mg(OH)2 solution

o Table 13. CO2 equilibrium data for 0.1 M Mg(OH)2 solution at 40 C

o Table 14. CO2 equilibrium data for 0.1 M Mg(OH)2 solution at 52 C

o Table 15. CO2 equilibrium data for 1 M Mg(OH)2 solution at 52 C

CHAPTER 1. LITERATURE SURVEY

1.1 Need for CO2 Capture and Sequestration.

The dependence on fossil fuels for achieving the primary energy needs of the world has

1 increased atmospheric emissions of CO2. The combustion of fossil fuels produces CO2, which is a known greenhouse gas believed to be responsible for nearly 60% of the global warming

effects. According to the IPCC 2001 report, CO2 contributes to half of the radiative forcing

2 which leads to the greenhouse effect. The percentage of CO2 in the atmosphere has increased by 96 % after the industrial revolution. 3 Fossil fuels are very likely to provide nearly 80% of the world’s energy requirements for the coming decades. The use of coal for electricity generation is rapidly increasing because coal is a relatively inexpensive source and electricity generation from coal accounts for 40% of greenhouse gas [GHG] emissions. 4

One of the first groups of emission sources being targeted for GHG reductions are coal- fired power plants because they are the largest stationary point source emitters of anthropogenic

5 CO2. CO2 emissions in the U.S. from coal combustion have increased over nearly 18% over

6 the period from 1990-2003 with forecasted 54% increase, if no CO2 control were to be applied.

If this increase is allowed to happen, the International Panel on Climate Change predicts that by the year 2100, the mean global temperature could have increased by around 1.9 C. This would lead to a sea level increase of 38 m, giving rise to species extinction and unmitigated disaster. 7

1.2 Various Technologies Involved in CO2 Capture and Sequestration

CO2 emissions are a cause of great concern to the world because of the implications to the

8 environment. There are three options to reduce CO2 emission into the atmosphere - a) to reduce energy intensity, b) to reduce carbon intensity and c) to enhance the sequestration of CO2.

The first option requires efficient use of energy. The second option requires switching to using non-fossil fuels such as hydrogen and renewable energy. The third option involves the development of technologies to capture and sequester more CO2.

Under the third option we have three technological pathways:

Air

Fuel Energy CO Post Combustion 2 Flue gas Conversion Separation

Power CO2

Air Air

H Pre Combustion Fuel Fuel CO 2 Energy 2 Power Conversion Separation Conversion

Power CO2 Flue gas

Fuel

O2 Oxy Combustion Air Air Energy Power Separation Conversion

N2 CO 2

9 Figure 1. Block flow diagrams of the three technological pathways for CO2 capture.

a) Pre-combustion capture

In this process, the CO2 separation takes place before the combustion of the fossil fuel.

The fossil fuel is converted into hydrogen by steam reforming or partial oxidation. 10 This is a

2 very common, well researched, industrial process used in the production of syngas and ammonia. 11 12 Partial oxidation is more advantageous when compared to reforming because it does not need an additional energy source.

13 Figure 2. Diagram of CO shift reaction in post combustion CO2 capture.

The heated H2O(g) and synthesis gas streams are fed into the CO-shift converter which

o 13 operates at 300 C and 2.14 MPa and CO2 is produced [CO+H2O -> CO2 + H2]. The CO2 can be captured by using a physical solvent [Selexol] which is far more energy efficient than other

14 physical and chemical solvents investigated. The CO2 separated would be sent to a compression unit and the hydrogen generated would be used as an input in a combined cycle to produce electricity. Several advancements are going on in these fields which include development of membrane-based gas separation systems to combine the gas shift reaction and hydrogen separation in one step. 15 16 This is part of the IGCC [Integrated Gasification

Combined Cycle] concept. The main advantage of this capture technique is that high partial pressure CO2 gas is generated, which results in availability of more driving force for separation.

However this system is limited by low availability, high investment and production costs. 17

3 b) Oxy-combustion

18 The concept of oxyfuel-combustion for CO2 emission capture was proposed in 1982.

The key concept in this process is that nitrogen is totally removed from the combustion air, by means of an air separation unit. A mixture of 95% pure oxygen and recycled flue gas are used.

The main reasons for recycling the flue gas are: a) A gas ready for sequestration is produced which consists mainly of CO2 and H2O b) Controlling the flame temperature

19 c) Making up the volume of missing N2 for carrying the heat over to the boiler.

CO2 has different properties when compared to N2 [denser than N2 and higher specific heat capacity]. So therefore flue gas recycle ratio and oxygen concentration during the combustion process become key parameters. 20

Air Air Separation N Unit 2

Oxygen

CO2 rich Concentrated Fuel Boiler or Gas flue gas Ash Removal/ CO2 Purification/ CO to storage Turbine Condenser/FGD Compression 2

Steam Recycled Flue gas

Steam Turbine Power

21 Figure 3. Oxyfuel technology for CO2 capture and storage.

Several investigations carried in pilot plant facilities worldwide have yielded theoretical and

22 23 design data for oxy-combustion CO2 capture. Very pure CO2 can be produced and this technology has a lot of capability for retrofitting with existing facilities.

Oxyfuel-combustion/sequestration is less efficient per unit of energy produced because it must provide power to several significant unit operations. 24 Large electric power demands must be met to operate cryogenic air separation units for generating pure oxygen.

Other problems include corrosion due to buildup of corrosive gases and the need to have high temperature resistant materials.

c) Post-combustion capture

NO ASH SULPHUR CO FLUE GAS COAL + AIR X 2 REMOVAL REMOVAL P-7 REMOVAL REMOVAL TO STACK

CO BOILER SELECTIVE ELECTROSTATIC FLUE GAS 2 CAPTURE UNIT CATALYTIC PRECIPITATOR DESULPHURIZATION REDUCTION UNIT

STEAM TURBINE FLY ASH GYPSUM CO2 STORAGE

Figure 4. Typical post combustion CO2 capture process.

The removal of CO2 from flue gas after the combustion of coal is referred to as post- combustion capture. Existing fossil fuel powered plants generate a flue gas that is at atmospheric pressure and which has a CO2 concentration of less than 15%. Therefore the driving force for

CO2 capture is low because the CO2 partial pressure is less than 0.15 atm.

Despite this problem, post-combustion CO2 capture has the greatest short-term potential for reducing CO2 emissions. The reasons are as follows:

1) There are around 5000 pulverized coal combustion power plants worldwide and 600 in the

United States of America alone. 25 At least for the coming decades, this will continue to be the type of power plant in operation. 26 Oxy-combustion and pre-combustion capture cannot be applied to these systems, but post combustion capture options can be easily retrofitted into the existing coal-fired power plant technology.

2) Flexibility of operation is available. Post-combustion capture operates as an independent unit and can be shut down in case of a failure. The other two capture processes are very integrated with the plant, so if capture fails, the entire plant must be shut down. 27

3) Post-combustion capture is the only available technology for gas-fired power plants.

4) Commercialization of IGCC technologies has fallen away with time and has not caught on in the way that it was expected to.

1.3 Challenges facing Post-Combustion CO2 Capture

Post-combustion CO2 capture has been used for many years in different industrial applications like natural gas treatment. 28 But several challenges have to be overcome in order to have a completely functioning, economic CO2 capture system to treat the flue gases for an entire plant. The additional energy requirements are as follows: a) Heat energy from the turbine required for operating the absorption and stripping process b) Energy requirements for compression, transportation and storage.

95% CO2 COMPRESSION LIQUEFACTION

TRANSPORT CO2 CAPTURE UNIT

STORAGE

ELECTRICAL ELECTRICAL COMBUSTOR/REHEAT ENERGY ENERGY BOILER REQUIREMENTS REQUIREMENTS

E-14 STEAM

HEAT ENERGY REQUIREMENTS

STEAM TURBINE

ELECTRIC POWER

29 Figure 5. Additional energy requirements for CO2 capture.

Two key factors which need to be considered before applying carbon capture and storage technology [CCS] to coal-fired power plants are the energy penalty and efficiency penalty. It is defined in the following two ways:29

EP = 100 x (Power output without CCS – Power output with CCS)/(Power output without CCS)

EP = Efficiency without CCS (%) – Efficiency with CCS (%)

For three leading CO2 capture techniques, the published values are given as follows:

Table 1. Energy & efficiency penalties for various CO2 capture techniques

CO2 Capture Method Energy Penalty Efficiency Penalty

Absorption by MEA 30 25 -27% 11-13%

Absorption by KS2 31 15 - 25 % 6 -10%

Oxy combustion 32 21% 9%

The values calculated above are mainly simulated values and for practical CO2 capture, the

29 energy penalties are much higher (around 40%). The largest application for CO2 capture is electricity production, but this energy penalty will reduce net electric output by 20-30% below

33 the plants which do not employ CO2 capture.

Costs for CO2 capture are reported as ($/tonne) CO2 captured or ($/tonne) CO2

34 avoided. As seen above, the CO2 capture processes are energy intensive and if implemented, will increase the cost of electricity for operating the coal-fired power plant. 35 The average cost for CO2 capture and compression has been worked out to be around $58/tonne of CO2 avoided

36 which results in a 63% increase in electricity cost for the power plant. The process of CO2 capture accounts for about 70% of total carbon capture and storage cost. This is the reason why most research efforts are directed towards this area. The need for a cost effective, potent, energy efficient CO2 capture technology is paramount in order to meet the Department of Energy’s CO2 mitigation technology goal by 2020.

1.4 Post-Combustion CO2 Capture Pathways

Post-combustion CO2 capture can be divided into five main categories:

POST COMBUSTION CO2 CAPTURE

ABSORPTION [CHEMICAL AND ADSORPTION MEMBRANES BIOLOGICAL CAPTURE OTHER NEW TECHNIQUES PHYSICAL]

Figure 6. Post Combustion capture pathways.

1.5 Post-Combustion CO2 Capture using Absorption

Absorption is a physical and/or chemical process in which gas molecules are dissolved in the liquid phase. Absorption is a common process in the chemical industry and is used to treat

37 industrial waste gas or flue gas streams containing acid gases like H2S, NOx, and CO2. An ideal solvent would have high reactivity towards CO2, low regeneration costs, high absorption

38 capacity, high thermal stability, low environmental impacts and low solvent costs.

1.5.1 Physical Absorption using Solvents

Physical absorption of CO2 using solvents takes place at low temperatures and high pressures. The amount of CO2 absorbed for these systems is directly proportional to its partial pressure following Henry’s law. The desorption can be achieved by increasing the temperature

(Temperature Swing Absorption) or lowering the pressure (Pressure Swing Absorption). 39 The

9 interaction between CO2 and the solvent is either weak Van der Waals type or electrostatic, which reduces the energy requirements for regeneration.

Physical absorption has been used commonly in the industry for CO2 removal from syngas in ammonia and methanol production. Some commonly used physical solvents are:

1.5.1.1 Sulfolane

Sulfolane solvent was first developed in the 1960’s by the Shell Oil company and it was used in natural gas purification. Sulfolane has many advantages such as high capacity, low vapor pressure (1.9 kPa at 150 °C), and low enthalpy of vaporization (543 kJ/kg ). 40 However, the most beneficial use of sulfolane in CO2 absorption is in the form of a mixed solvent with other amines such as Diisopropyl amine (DIPA), 2-amino-2-methyl-1-propanol or methyl diethanolamine (MDEA). These mixed solvents combine the advantages of physical and chemical solvents such as increased absorption capacity (reducing stoichiometric dependence),

41 higher purity of treated CO2 and lesser recirculation of the solvent. The solvent with sulfolane and DIPA is referred to as Sulfinol D and the solvent with Sulfolane and MDEA is called

Sulfinol M. This solvent has many advantages including low solvent recirculation rates and low solvent degradation rates. The operation of a Sulfinol unit is similar to that for an alkanolamine solution but for the addition of a flash tank. The Sulfinol process operates at maximum

2 42 efficiency at a H2S/CO2 ratio of 1:1 and an acid gas of pressure 110 lb/in . abs. The flash tank is required to dispose the hydrocarbons that have been absorbed in the solvent and prevent operational difficulties. 28 The main disadvantages with this technique are high hydrocarbon solubility and the high cost of chemicals used in the process.

1.5.1.2 Selexol

The Selexol process uses a physical solvent to remove acid gas from streams of synthetic or natural gas. This solvent has been in use from the 1960’s and is a mixture of dimethyl esters

43 of polyethylene glycol [CH3 (CH2CH2O)n CH3], where n is between 3 and 49. This system is highly suited for the bulk removal of H2S and CO2. The key driving force for the Selexol process is the acid gas partial pressure. The absorption takes place at low temperatures

(0–5 °C) and desorption takes place by pressure reduction or stripping with air or steam. 44

Selexol is best suited for applications with high CO2 content in the acid gas stream. The Selexol solvent is also chemically and thermally stable and it suffers from minimal losses due to its low vapor pressure. This solvent also has preferential solubility for H2S and therefore it is ideally suited for H2S removal in the presence of high CO2 content. The limitations of this technique include dependence of efficiency on high pressure and potential hydrocarbon losses due to high affinity of Selexol towards hydrocarbons.

1.5.1.3 Rectisol

The Rectisol process was one of the earliest used physical solvents for acid gas treatment and has been used for synthesis gas applications. Chilled methanol is used as the solvent and is very complex compared to other physical solvent processes. The high vapor pressure of methanol necessitates the operation of the process at temperatures ranging from 0 oC to -70 oC. 45

The solubility’s of H2S and CO2 are high in methanol compared to Selexol and also allow for deep sulphur removal. Regeneration can be easily achieved by flashing the rich solvent at low pressure. Other advantages of this process include removal of carbonyl based compounds and high miscibility of the solvent with water, eliminating foaming losses. The need to refrigerate

46 the solvent and the complex operating scheme are the main disadvantages of this process.

These disadvantages can be offset by the net reduction in solvent flow rate for CO2 removal compared to other physical solvent processes. 47

1.5.1.4 Fluor Process

The Fluor Process uses propylene carbonate as a solvent and is used generally in syngas treatment. It is licensed by Fluor Daniel, Inc. and is particularly useful for treating acid gas steams with high CO2 partial pressure. Propylene Carbonate (C4H6O3) is a polar solvent which has higher CO2 solubility compared to any other physical solvent. A typical Fluor process has two stages of flashing and is ideal suited for acid gas streams with an acid gas partial pressure

> 60 psi. 48 The operating temperature for the Fluor solvent ranges between -18 oC to 65 oC.

Despite the vapor pressure of propylene carbonate being higher than Selexol, it has much lesser solvent losses. The Fluor process requires no water wash to recover the solvent because of its low vapor pressure. 45 The major drawbacks of this process include need for high solvent recirculation, expensive nature of the solvent and unstable nature of the solvent when H2S is present in more than trace concentrations. 44 45

1.5.2 Chemical Absorption using Solvents

Chemical absorption systems are the most preferred option for post-combustion CO2 capture because they are ideally suited for treating flue gases with low to moderate CO2 partial pressure. Absorption using chemical solvents at this point in time offers the highest capture efficiency and lowest energy costs when compared to other existing post-combustion capture systems. 49 The flue gas is cooled and brought into contact with the solvent in an absorber/contactor. The absorption takes place due to a chemical reaction between the CO2 and the solvent. The absorber temperatures are between 40 °C and 60 °C and the CO2 binds weakly

12 with the chemical solvent to form an intermediate compound. The CO2 rich solution is pumped to a stripper vessel and the solvent is regenerated by using high temperatures (100 °C – 140 °C) and pressures.

1.5.2.1 Amine-Based Solvents

One of the most commonly used and studied solvents are the amine-based solvents, particularly monoethanolamine [MEA]. It was developed over 60 years ago as a general, non-

50 selective solvent to remove acidic gas impurities (e.g. H2S, CO2) from natural gas streams.

Pure unhindered amine reacts with CO2 to form a water-stable compound called carbamate and

51 half a mole of CO2 is absorbed per mole of amine. A conventional MEA would involve cooling the flue gas and contacting it with a lean solvent with a CO2 loading of 0.1-0.2 mol

CO2/mol MEA, yielding a rich solvent of about 0.4-0.5 mol CO2/mol MEA loading at a temperature of 40-60 °C. 52 The rich solvent is pumped to a stripper, where it is regenerated at elevated temperatures (100-120 °C) and pressures (1.5-2 atm). 53 53a Major problems associated with amines are: a) solvent degradation, b) high energy consumption for regeneration [1900

54 51 kJ/kg CO2], and c) corrosion.

Several ethanolamine derivatives can be produced by replacing one or more of the alcohol groups by hydrocarbon groups e.g. methyldiethanolamine (MDEA). MEA can react more than MDEA, but MDEA has a greater CO2 absorption capacity [1 mol of CO2/mol amine]

55 56 57 and requires lower regeneration energy [1200 kJ/kg of CO2]. Mixed amines and amine blends have been reported to maximize the desirable qualities of the individual amines. 58 There could be reduction in cost for regeneration of the solvent by up to 20%, if a mixture of MEA and

MDEA is used. 59 This is due to the lower heat of vaporization of MDEA [223 cal/g] when compared to MEA [355 cal/g]. It has also been shown that the addition of a small amount of

MEA to an aqueous solution of MDEA significantly increases the enhancement factor and the

60 rate of CO2 absorption.

1.5.2.2 Carbonate-Based Solvents

Aqueous-based carbonate solvents comprise the second most popular class of solvents used in acid gas treating. A soluble carbonate reacts with CO2 to form bicarbonate, which when heated, releases CO2 and the solvent is regenerated as a carbonate. A substantial reduction in the energy requirement, compared to amines, is the major advantage of this process. However, their low reaction rates require a use of a promoter like piperazine to increase the rate of reaction. 61.

Piperazine has a cyclic diamine structure that favors rapid formation of carbamates and it also

62 catalyzes proton extractions in the reaction mechanism. The amine can absorb 2 moles of CO2 for every mole of amine and with potassium carbonate, which has capacity of 1 mol CO2/mol

63 K2CO3 in solution, the blended solvent has potential for high CO2 capacity. When piperazine is blended with K2CO3, the amount of amine protonation is reduced by the buffering capacity of

64 the potassium bicarbonate, which leaves more amine free to react with CO2. The K2CO3/PZ system (5 molar K; 2.5 molar PZ) has an absorption rate 10–30% faster than a 30% solution of

MEA. 61 A model was simulated in Aspen plus in which a five-stage absorption-desorption unit was used. Five stages were chosen during modeling because it gave a good compromise between calculation precision, computational effort and convergence behavior of the model. 65

The absorber was operated at 40 °C and stripper was also operated at nearly the same temperature. This is possible because the desired heat of absorption of the promoted potassium carbonate is near or below that of water. 63

Pure piperazine itself has been investigated as a novel solvent for CO2 absorption.

Piperazine has was thought to have limited solubility, but recent research shows that PZ has the

66 capability to be used in very high concentrations when it is partially loaded with CO2. The

CO2 absorption rate is twice that of 30% MEA and the thermal degradation also is negligible up to 150 °C. 67 PZ will also require 2 to 3 times less absorber packing than MEA. 68 Piperazine is

3 to 10 times more expensive than MEA, but this cost can be offset by other advantages of piperazine, namely its resistance to oxidative degradation and its corrosion resistance to stainless steel. 33

1.5.2.3 Ammonia

Ammonia based wet scrubbing systems are very similar in working to the amine based systems. It reacts with CO2 mostly in the form of ammonium carbonate (AC) and ammonium bicarbonate which has lower heats of reactions than amine based systems. An additional advantage for this system is that when it reacts with SO2 and NOx, fertilizers are formed which is a salable by product. 61 A typical removal efficiency can reach 90% with 5% or higher concentration, but if the concentration exceeds 15%, substantial amount of ammonia will volatilize from the solution. 38 The major problems associated with this process are that the flue gas must be cooled to ambient temperature to increase the CO2 uptake and to reduce the ammonia vapor emissions from the absorption column.

Another novel process which involves the use of ammonia is the chilled ammonia process, where the absorption takes place at a very low temperature of 0-10 °C which prevents the ammonia

69 from evaporating. In this process, the CO2 loading concentration of the CO2 lean stream ranges from 0.33-0.67 mol CO2/mol NH3 and CO2 rich stream from 0.5-1 mol CO2/mol NH3.

The desorber temperature ranges from 50-200 °C and the pressure from 2-136 atm. 70 The energy required for this process [1147 kJ/kg CO2] is much lower than the corresponding values

71 for MEA and amines[4215 kJ/kg CO2].

1.5.2.4 Dual Alkalis

The Solvay process makes use of two alkalis to sequester CO2 in the form of sodium carbonate which is a very useful by-product. Since the Solvay process is not efficient, a

72 modified dual alkali approach was proposed. This process involves the reaction of CO2 with

NaCl and NH3 to produce sodium carbonate. The primary alkali ammonia is recovered by reacting the NH4Cl produced, with lime. Lime acts as the secondary alkali and limestone is used as a source of lime. The drawbacks with this system include consumption of limestone, intensive energy requirements and also the release of one mole of CO2 for every mole of CO2 captured during the calcination of limestone.

1.5.2.5 Ionic Liquids

Ionic liquids are a broad category of salts, typically containing an organic cation and either an inorganic or organic anion. Ionic liquids are non-flammable, environmentally benign and can exhibit exceptional thermal stability. 9 They are known as designer solvents because their physical properties such as melting point, viscosity and gas solubilization can be controlled by altering the substituents of the cation or anion. 73 The mechanism for capture is based on physisorption and involves weak interaction between the ionic liquids and CO2 with a heat of absorption around -11kJ/gmol. ILs can dissolve gaseous CO2 and are stable at temperatures up to several hundred degrees centigrade. Their good temperature stability offers the possibility of recovering CO2 from ﬂue gas without having to cool it ﬁrst. Since ionic liquids are physical

16 solvents, little heat is required for regeneration. However ionic liquids face problems like corrosion, expensive bulk scale production and very high viscosities [66 to 1110 cp at 25 °C]. 74

1.5.3 Absorption using Carbonation/Calcination Cycles

This capture technique involves the reaction of CO2 with a solid metal oxide yielding a metal carbonate as the product and regeneration of the metal oxide by heating the metal oxide beyond the calcination temperature. This system has several advantages like: a) CO2 separation can be performed at flue gas conditions without the need for high system pressure/low temperature. b) High equilibrium capacities of the sorbent (393 g of CO2/kg) compared to the amine process (MEA- 60 g of CO2/kg). c) Ability of generate pure streams of CO2 despite the

75 presence of SO2 and other gases. Calcium oxide has proposed as the sorbent of choice for this method because of its high absorption capacity and cyclic stability. The cyclic CCR

(carbonation calcination reaction) process has two reactors: 1) carbonation reactor – where CaO is carbonated to CaCO3 at 600 °C -700 °C and atmospheric pressure. 2) calcination reactor – where calcination of CaCO3 takes place to regenerate the sorbent and produce a concentrated

76 stream of CO2 at high temperatures (> 900 °C). Fluidized beds of CaO have been shown to

77 give excellent CO2 capture efficiencies from combustion flue gases at high temperatures.

Lu et al. 78 demonstrated the applicability of this technology using a pilot scale dual fluidized system, with continuous sorbent looping and achieved high capture efficiency (> 90%) for the first several cycles. The major disadvantage of this technique is the decrease in absorbent activity with the increase in the number of carbonation/calcination cycles. This takes place due to various factors like sintering, sulphation and attrition. Some potential methods for reducing problems caused by sintering include spent sorbent reactivation by air/steam and thermal pretreatment of the sorbent. 79

1.6 Post-Combustion CO2 capture using Adsorption

Adsorption is a heterogeneous process which involves the capture of CO2 molecules on the surface of a solid sorbent. The CO2 interacts with the sorbent via weak Vander waals forces

(physisorption) or stronger covalent bonds (chemisorption). The advantages of using such sorbents for CO2 capture include reduced energy costs for regeneration and greater capacity and selectivity. Regeneration of the solid sorbent is done via Pressure Swing Adsorption (PSA) or

Temperature Swing Adsorption (TSA). Pressure Swing Adsorption is a cyclic process where the adsorption is carried out a relatively higher pressure compared to the desorption pressure, with a part of the product from the adsorption step. 80 PSA has been used more extensively due to lower energy demand and higher regeneration rate.

Activated carbon based solid sorbents and carbon molecular sieves have been studied extensively for flue gas CO2 capture. Activated carbon based systems operate at close to room temperature and atmospheric pressure. Initial studies indicated that around 70% CO2 removal was achievable using AC’s at 298 K coupled with PSA. 81 Siriwardane et al. 82 demonstrated that activated carbons at higher pressures (>25 psi) show significantly higher CO2 capacity than

16 molecular sieves. Garcia et al. observed that increase in temperature leads to reduction in CO2 capacity and breakthrough time for the activated carbon sorbent. Activated carbons can be produced from a wide variety of sources and thus can be produced at a lower cost on the industrial scale. They also possess other significant advantages such as hydrophobicity and lower energy requirements to carry out regeneration [20 kJ/gmol]. 83 84 But, these advantages are offset by other problems like low absorption capacity at lower partial pressures of CO2, reduction in capacity with temperature and low selectivity.

Carbon Molecular Sieves (CMS) are a class of microporous carbon materials that are used for separating molecules on the basis of molecular weight or size. These CMC’s must exhibit a highly developed pore network and controlled size opening to differentiate between molecules having similar dimensions. Initial studies performed by Kapoor et al. 85 demonstrated the applicability of carbon molecular sieves with PSA for separation of a CH4/CO2 mixture.

Burchell et al. 86 developed a porous monolithic activated carbon material which has a large

87 micropore volume and achieved a CO2 uptake of >100 mg/g. Wahby et al. prepared a series of CMS’s from petroleum pitch using potassium hydroxide as an activating agent and achieved high capacities of CO2 absorption (380 mg CO2/g sorbent) at 273 K and 1 bar. The major drawbacks of this technique include the need for pretreatment of gases before adsorption, low

CO2/N2 selectivity and temperature dependence on adsorption capacity.

1.7 Post-Combustion CO2 Capture using Membranes

Membranes are semi permeable barriers which can selectively remove CO2 from flue gas through mechanisms like diffusion, ionic transport and adsorption. Membranes can either be constructed by using either polymeric or inorganic material and can be porous or non porous. 44

Membrane based processes have generated a lot of interest because it can significantly bring down the cost for regeneration of the sorbent. Membranes have not yet been commercially employed for CO2 capture from flue gas, but have been used extensively for CO2 recovery from natural gas. 88 Three major variables that play an important role in membrane based separation processes are:

a) Membrane selectivity (α) – intrinsic parameter determined by the ratio of pure gas permeabilities, b) Pressure ratio (ψ) – is the ratio of pressures between upstream and downstream compartments, c) Stage cut (θ) – is the ratio of the permeate flow to the feed flow rate.89

Membrane process for CO2 separation can be divided into three main systems: a) non dispersive contact via microporous membranes, b) Gas permeation systems c) Supported liquid membranes. 90 Some of the commonly used membranes in these systems are discussed below:

1.7.1 Polymeric Membranes

Polymeric membranes are classified into rubbery, polyolefinic or glassy depending on their operating temperature relative to the glass transition temperature. 91 The major mechanism in the transport of gas molecules through a polymeric membrane is by solution diffusion The key relation between permeability, diffusivity and solubility for a polymeric membrane is given by P = DS, where P is the permeability coefficient – a measure of the flux of the membrane (cm3 cm-2 s-1 cm Hg-1), D is the diffusivity coefficient (cm2 s-1) and S is the solubility coefficient which measures the solubility of the gas molecules within the membrane (cm3 cm Hg-1). 92

Achieving a higher permeability and selectivity for CO2 is the main focus of membrane research and development. Polymeric membranes cannot withstand temperatures above 100 ° C and also experience a drop in selectivity with increasing temperatures. So therefore, polymeric membranes are more suited for post-combustion because the flue gas temperature is low. But lower partial pressures of CO2 necessitate the need for compressors or vacuum systems, which increase the overall cost of the CO2 capture process. Single stage membrane systems would result in low efficiency, as the permeate would be diluted with nitrogen gas, thereby creating the need for an additional membrane step. 93 Polyamide membranes with the 6FDA group have been

94 researched extensively and give good selectivity for CO2. Polyethylene oxides (PEO) membranes have also been found to be good candidates for CO2 absorption owing to strong

95 affinity for the CO2 molecule by the polar ether oxygen in the polymer chain. The major drawback identified for the PEO membranes is its tendency to crystallize, which causes a

20 decrease in permeability. Other technologies investigated for improving the performance of the membrane include mixed matrix membranes (incorporation of an inorganic nanoparticle into a polymeric matrix) and facilitated transport membranes (coating the membrane with cross linked polyvinilamine). 96

1.7.2 Zeolitic Membranes

Zeolites are porous, inorganic, crystalline aluminosilicates that have been used extensively to prepare membranes to separate CO2 from CH4. Initial studies revealed

97 physisorption to be the dominant process for CO2 absorption on zeolites. The uniform crystalline nature of zeolites with well-defined pore sizes between 0.5 to 1.2 nm also allows for separation of molecules through the molecular sieving effect. Zeolites have excellent regeneration properties compared to other membranes/sorbents and possess extremely fast adsorption kinetics. Both natural and synthetic zeolites have shown promising results for CO2

98 separation from flue gas streams. Siriwardane et al. performed CO2 absorption studies on five zeolites (4A, 5A, 13 X, APG-II and WE-G 592) at 120 °C and concluded that zeolites 13 X and

WE-G 592 showed the best CO2 absorption capacities (0.7 mol/kg at 1 atm and 1.2 mol/kg at 20 atm). They also concluded that absorption capacity is proportional to the pore diameter and the

Na/Al ratio of the zeolites. Further studies showed that zeolite 13X is extremely selective to

99 CO2, possessing high isosteric heats of adsorption (37.2 kJ/mol). Temperature swing adsorption has been investigated by many research groups for regeneration of the zeolites and zeolites 5A/13X show excellent CO2 uptake capacity combined with good regeneration capability. However zeolites suffer from a lot of disadvantages including attrition of the zeolites, detrimental effect of water on CO2 adsorption and decrease in adsorption capacity with temperature.

1.7.3 Membranes in Conjunction with Chemical Solvents

These hybrid membrane-solvent systems are being researched actively because they offer the possibility of achieving high selectivity and high permeabilities. The solvent is supported on the pores of the membrane or introduced inside the pores of the membrane. They provide a high surface area to volume ratio and ensure excellent mass transfer between the gas and the solvent stream. The immobilization of aqueous amine on to a solid support like silica has several advantages like bringing down the energy required for regeneration, reducing the degradation of the amine and cutting down on the corrosion of the absorber. Supported amine adsorbents can be divided into three classes: a) Class 1: porous supports impregnated physically with monomeric or polymeric amines, b) Class 2: amines covalently linked to the supports, c) Class 3: amines which are polymerized in situ from an amine monomer. 100 The most commonly used amine for impregnation is Polyethylenimine (PEI), which is a branched amino polymer or a linear amino polymer with secondary amines. The effect of PEI impregnation was studied on various carbonaceous materials like activated carbons, fly ash, biochar and carbon nanotubes and a substantial increase in CO2 adsorbate interactions were found. However, these methods could result in potential blocking of the pores, whose mechanisms have not yet been fully understood

Xu et al. 101 developed a novel molecular basket adsorbent (Class 1) by dispersing PEI into pores of mesoporous MCM -41 and selectively achieved CO2 adsorption from flue gas, maintaining the stability of the sorbent after several cycles. PEI-impregnated MCM-41 showed an increase in adsorption capacity with increasing temperature compared to other conventional adsorbents like activated carbons and zeolites. 102 Several other silica supported sorbents have been studied (SBA-15, MCM-48, PE-MCM-41 and SBA-16) and all of these class 1 supported amines show excellent CO2 sorption capacities compared to the pure amines. Although Class 2

supported amines (amine grafted covalently on to the silica support) show improved CO2 - adsorption capacities, they suffer from problems like low thermal stability in the presence of water vapor and lower equilibrium CO2 adsorption capacities compared to class 1 supported amines. Class 1 supported amines showed improved performance in the presence of water vapor

103 and can effectively adsorb CO2 with high working capacities of 4 mmol/g. Supported amines show great potential in bringing down the energy requirements for the CO2 capture process, but are very much in the nascent stage of development and require further investigation about the selection criteria and the effect of other gases on these sorbents.

1.8 Post-Combustion CO2 Capture using Biological Methods

500 billion tons of CO2 are fixed annually by terrestrial vegetation and microalgae have

104 the ability to fix CO2 with around 10 times efficiency as that of terrestrial plants. A controlled photosynthesis reaction by the microalgae that fixes CO2 has been investigated as it has additional advantages of high value biomolecule production, production of H2 and reduction in the cost of CO2 capture. Several studies have shown that microalgae are capable of growing rapidly under high CO2 concentrations and also in the presence of SOx and NOx , which are common components of flue gas. 105 106 The microalgal strain Chlorella Vulgaris shows excellent growth at high temperatures and CO2 concentrations of 15% (which is the average flue

107 gas CO2 partial pressure). A photo bioreactor with Chlorella Vulgaris exposed to 1850 ppm

108 of CO2 and short gas residence was able to remove upto 74% of the CO2 in the airstream.

However, CO2 mitigation by microalgae would be economically feasible, only when the biomass produced leads to products that have substantial commercial value. Another major disadvantage faced by these microalgae systems are that they require land and water resources in close proximity to the power plant.

1.9 Other Technologies for CO2 Capture

1.9.1 Cryogenics

Cryogenic separation is a novel technique where all other components of the flue gas are removed barring N2 and CO2 are removed and the remaining gas is sent into a chamber where

CO2 is liquefied by fractional condensation and distillation at low temperatures. When the system operates at the triple point of CO2 (-56.6 °C and 7.4 atm), the CO2 will liquefy into a concentrated solution and the N2 can be vented out at the top of the chamber. The major advantage of this process is that the liquid CO2 can easily be transported, stored or sent to enhanced oil recovery fields. 109 Tuinier et al. 110 demonstrated the applicability of cryogenic capture to post-combustion CO2 by achieving effective separation of CO2/N2 mixtures using dynamically operated packed bed systems. The Cryocell process developed by Cool energy Ltd showed excellent CO2 capture efficiencies from natural gas, comparable to amine based technology and had many advantages like the non-requirement of chemicals, water supply and heating systems. 111 The potential roadblocks with this technology are the high cost of energy for refrigeration and possible plugging by ice in the feed streams of the cooling units.

1.9.2 Metal Organic Frameworks

Metal organic frameworks (MOF) are a new class of crystalline porous material comprised of metal containing nodes linked by organic ligand bridges and held together by coordination bonds. 112 The MOF materials have 3D structures with uniform pores and a network of channels that are stable to the elimination of guest molecules (solvents) and can store gases like H2 and CO2. The pore size, topology and chemical composition of the MOF’s can be tuned over a wide range resulting in the maximum reported capacities for CO2 adsorption.

Additionally, MOF’s possess fast adsorption rates which are essential for gas separation.

113 Yazaydin et al. screened over 14 MOF’s for CO2 adsorption at flue gas pressures and reported that Mg/DOBDC and Ni/DOBDC show excellent CO2 capacities at 0.1 atm and 298 K.

Liu et al. 114 compared the performance of Ni/DOBDC and reported that in addition to better performance, Ni/DOBDC can resist the impact of water on adsorption better than zeolites. The major problems to be overcome in this area include high synthesis cost of the MOF’s and improving the stability of the MOF’s towards heat regeneration, water vapor and acid gases.115

1.10 Conclusion

This chapter provided a concise overview of the available technologies for post-combustion CO2 capture. There is an urgent need to develop technologies and process that are able to achieve DOE’s goal of 90% CO2 capture with less than 35% increase in cost of electricity. Post-combustion capture using chemical solvents offer the best near term potential for CO2 capture due to their state of development and commercial viability for integration with power plants. The implications, advantages and process conditions for using magnesium hydroxide solutions as a post combustion solvent will be discussed in the upcoming chapter.

CHAPTER 2. MAGNESIUM HYDROXIDE AS POST-COMBUSTION

SOLVENT FOR CO2 ABSORPTION

2.1 Magnesium Hydroxide for CO2 Capture

Amine based solutions are the most widely used and studied systems for post combustion

CO2 Capture. A common solvent for reference is 30% weight aqueous MEA

[Monoethanolamine], which have been used previously by oil and chemical industries for acid gas removal from natural gas streams. But MEA suffers from several disadvantages which include: a) solvent degradation due to side reactions with SO2 and other flue gases, b) corrosion

116 of the equipment, c) low CO2 loading capacity and d) high regeneration energy. The requirement for a solvent that is non-corrosive, SO2 resistant and has a low binding energy for easy regeneration is paramount.

Magnesium based minerals like serpentine, brucite and olivine have been investigated for permanent CO2 disposal in the form of magnesite, which is an environmentally stable benign carbonate. It can also be used in the form of a slurry solution to absorb CO2 in a packed/bubble column and produce a rich stream of CO2 by recovering the solvent in a regenerator. The idea of using magnesium hydroxide solutions as a potential solvent to scrub CO2 from flue gases was derived from the established practice of Flue Gas Desulphurization using lime or limestone solutions. Magnesium hydroxide was found to give superior continuous absorption capacity

117 compared to other alkaline solutions like NaOH and Ca(OH)2. The process investigated in

ο this study involves CO2 absorption using magnesium hydroxide slurry solutions at 52 C and regeneration of the solvent at a stripper temperature of 65 οC.

2.2 Process Description

A process flow diagram was developed as shown below in the figure 7. The system requirements for using Mg(OH)2 as a solvent to separate CO2 from flue gas and sequester it with a 95%(mole) purity were estimated based on the molar compositions of 15% CO2, 13% H2O and

3% O2 and the remaining N2 at the inlet of the CO2 absorber with 90% CO2 removal from a typical 500 MW coal fired power plant. 4M Mg(OH)2 (21% wt of Mg(OH)2) was assumed to be used to achieve 90% removal.

Figure 7. Process flow diagram of the proposed setup.

Here, the flue gas is assumed to enter an absorber section at an adiabatic saturation temperature of 52 °C from the FGD system. The gas stream is contacted with the scrubbing liquor containing the magnesium compounds at a pH which is low enough to allow as much as possible of the magnesium carbonate compounds to be in solution, but high enough to allow for rapid

27 absorption of CO2. The rich magnesium slurry containing 1 % (mole) CO2 is raised from 52 °C to 62 °C after passing through a major heat exchanger (HX1) and further increased by 3 °C from

62 to 65 °C under 5.1 atm with low pressure steam in the stripper.

The operating pressure of the stripper was determined using the Henrys law data that take into account the dissociation of CO2 in the aqueous phase using PRO II process simulator. The operating pressure of the stripper is created by a pump (P1) and is maintained to have a purity of

95 %( mole) in the CO2 exiting the stripper. The high operating pressure of 5.1 atm is useful to reduce the amount of water vapor in the CO2 stream which in turn reduces the load on the LP steam and no additional cooling water is also required to meet the 95 % CO2 specifications. Also the subsequent downstream compression and dehydration processes would require much lesser energy.

2.3 Advantages of Using Magnesium Hydroxide for CO2 Capture

As a by-product of magnesium-enhanced FGD systems, the material will be readily available, and its production will not result in the release of additional CO2 emissions to the atmosphere. The active ingredient Mg(OH)2 is readily available commercially and can be reclaimed from any lime based FGD systems as all limestone’s have magnesium hydroxide up to some extent, ranging from 4 % to 45 %. The outlet temperatures of most flue gases are around

52 °C. If Mg(OH)2 is used for CO2 absorption, no additional heating of the flue gas is required.

Regeneration of Mg(OH)2 takes place at 65 °C in the stripper which implies due to the small temperature difference, lesser energy will be consumed during the stripper operation. CO2 separation is high with recycle and it has a very favorable chemistry for absorption (ΔG = -9.13 kcal/mol at 25 °C). 117 The scrubber can also react with other acid gases like sulphur oxides

(SO2 and SO3) coming from the FGD scrubber and can separate sulphur in the form of

CaSO4 .No modification of the existing boiler-heat transfer configuration is required and the system can be easily retrofitted into the existing FGD units or designed as new units in combination with the FGD unit.

2.4 Magnesium Hydroxide Recovery from Magnesium-Enhanced Wet FGD

Low cost magnesium hydroxide for CO2 capture can be generated from magnesium enhanced wet Flue Gas Desulphurization units originally developed by Carmeuse North

America. A magnesium-enhanced process using lime has around 3-6% of MgO in the calcium oxide. The MgO in the lime produces MgSO3 in the absorber liquid. MgSO3 is a soluble salt that accumulates in the liquid portion of the slurry and acts as a catalyst, by reducing the resistance to the transfer of SO2 from the gas to the liquid. The slurry has the various reaction products with SO2, liquid phase SO2 and solid calcium hemihydrate. Slurry is pumped from the reaction tank and sprayed into the flue gas. SO2 reacts with MgSO3 to form Mg(HSO3)2. Slaked lime consisting of Ca(OH)2 and Mg(OH)2 react with magnesium bisulfite to get calcium sulfite hemihydrate (CaSO31/2H2O) and MgSO3 is regenerated which is then available for continued capture of SO2 when reaction tank slurry is recirculated.

The reactions involved are

Absorption:

SO2 + H2O + MgSO3 -> Mg(HSO3)2

Precipitation and regeneration:

Mg(HSO3)2 + Ca(OH)2 - > CaSO3. ½ H2O (solid) + MgSO3 + 1-1/2 H2O

Mg(HSO3)2 + Mg(OH)2 -> 2 MgSO3

Mg(OH)2 production:

MgSO4 + Ca(OH)2 + 2 H2O -> CaSO4.2 H2O(gypsum) + Mg(OH)2

One portion of the absorber slurry containing the soluble magnesium sulfite and the insoluble calcium sulfite hemihydrate is sent to an oxidation tower and is further oxidized using compressed air at the bottom of the tower. Calcium sulfite dihydrate is converted to gypsum and magnesium sulfite is converted to magnesium sulfate. Magnesium hydroxide is produced by adding slaked lime to the liquid containing MgSO4. MgSO4 undergoes a double decomposition reaction to precipitate as magnesium hydroxide and gypsum.

Figure 8. Schematic of magnesium hydroxide recovery process. 118

2.5 Chemistry of the Absorption / Desorption Process.

* Carbon dioxide hydrolyzes upon dissolution in water. The product H2CO3 can be considered as the effective carbonic acid value in water due to difficulty in differentiating between dissolved CO2 and carbonic acid. At a pH of 8.4 – 8.6, the formation of the carbonate ion is minimal.

The reactions of interest for magnesium hydroxide solutions exposed to atmospheres containing

119 CO2 are:

* CO2 + H2O <-> H2CO3

* + - H2CO3 <-> H + HCO3

- + 2- HCO3 <-> H + CO3

2+ - Mg(OH)2 <-> Mg + 2OH

2+ 2- Mg + CO3 -> MgCO3

2+ - Mg + 2HCO3 <-> Mg(HCO3)2

Mg(HCO3)2 <-> Mg(OH)2 + CO2

The magnesium bicarbonate formed is completely soluble and does not exist as a solid.

The solid magnesium carbonate (MgCO3) formation is controlled kinetically and requires more time than the liquid residence time of the system. So for successful operation, the system needs to be operated at a high pH which will allow CO2 absorption, but low enough to nullify the formation of magnesium carbonate.

The objective of this thesis is to investigate three crucial phenomena in this absorption process, namely:

1) The rate of dissolution of magnesium hydroxide and impact of factors like pH,

temperature, stirring rate and particle size.

2) Kinetics of magnesium carbonate formation and the impact of pH control and

temperature.

3) Vapor-liquid-solid equilibrium of the CO2-H2O-Mg(OH)2 system.

CHAPTER 3. DISSOLUTION RATE STUDIES OF MAGNESIUM

HYDROXIDE

3.1 Introduction

One of the most important steps in the CO2 scrubbing process is the rate at which the magnesium hydroxide dissolves from the solid part of the slurry and becomes available for reaction with the absorbed CO2. A thorough understanding and prediction of the dissolution rate process under different conditions is required for the successful operation of the CO2 absorption system.

120 CO2 gas absorption into Mg(OH)2 solution takes place based on the following steps.

1. Diffusion of CO2 gas through the gas film near the gas liquid interface

2. Dissolution of CO2 in the liquid phase.

3. The first dissociation of CO2 [into bicarbonate ions].

4. The second dissociation of CO2 [into carbonate ions].

5. The dissolution of Mg(OH)2 from the solid part of the slurry

6. The diffusion and subsequent reaction of inorganic carbon species within the reaction

zones of the liquid film.

Step 1 depends on the type of equipment used to bring the gas into contact with the slurry solution. Step 2 – 4 take place instantaneously. Magnesium hydroxide is nearly insoluble in water and has a Ksp value of 5.61 *10-11 (mol3/dm9). So therefore, the dissolution rate of

Mg(OH)2 is the slowest step in this reaction and therefore very likely to be the rate determining

33 step of the entire reaction. Thus prediction of the dissolution rate of magnesium hydroxide is essential for the design of the Mg(OH)2 based - CO2 absorption system.

Over the years, several studies have been conducted on the dissolution of limestone

(calcium carbonate), which is used as a sorbent for SO2 capture in wet Flue Gas

Desulphurization (FGD) systems. Most of the earlier studies focused on the impact of dissolution on geological processes such as evolution of Karst landscapes in limestone areas 121

122, effect of calcite dissolution on ocean chemistry 123 124 125 , growth and dissolution of sedimentary rocks 126 127 and acid fracturing of oil field reservoirs 128 129. Plummer et al. 130 conducted a critical review of the existing work and concluded that dissolution rates were controlled by hydrogen ion diffusion below pH 5 and by surface reaction kinetics above pH 5.

More recent studies have focused on estimating dissolution of limestone, lime and other calcium based materials under conditions encountered in wet FGD units.

131 Chan and Rochelle measured the dissolution rate of reagent CaCO3 at constant pH using batch titration with HCl and developed a mass transfer model assuming that the calcite particles behaved as spheres in an infinite stagnant solution. Toprac and Rochelle 132 showed that there is no influence of the type of limestone on the dissolution rate and that dissolution rate depends heavily on the particle size distribution.

Gage and Rochelle 133 studied the impact of sulfite on dissolution and concluded that, in the presence of sulfite, the dissolution rate was controlled by a combined mass transfer/ surface reaction regime. Tseng and Rochelle 134 also investigated the dissolution rate of calcium sulfite

(by product of FGD systems) and showed that under simulated FGD conditions, the dissolution rate solely depends on mass transfer.

Other studies have also discovered that under certain conditions, surface reaction may also influence the dissolution of limestone. Lund and Fogler 135 studied limestone dissolution in

HCl using a rotating disk apparatus and found that at -15.6 °C both mass transfer and surface reaction controls the dissolution rate. Gao Xiang et al. 136 measured limestone dissolution rate under conditions of pH (4 – 6), temperature (25°C -55 °C) and dissolved sulfite (1mM). They obtained the same trend of results reported by Gage and Rochelle and also concluded that the dissolution rate of limestone is inversely proportional to its crystallinity.

Siagi et al. 137 investigated the dissolution rate of South African calcium based material using HCl solution to maintain a constant pH and concluded that the dissolution rate was controlled by surface chemical reaction control. Their studies showed that the dissolution rate increased with an increase in temperature and decrease in particle size. The dissolution kinetics of magnesium oxide during the leaching of magnesite was found to be controlled by surface chemical reaction with high values of apparent activation energy (58-64 kJ/gmol) 138 The dissolution of calcined magnesite was also reported to follow chemical reaction control, which was further supported by a linear relation between the particle size and the apparent rate constant . 139

Various studies have also searched for ways to increase the rate of limestone dissolution by using various additives. It was also reported that organic acids with buffering capacity between the pH of the gas/liquid interphase and pH of the bulk liquid phase would provide optimum enhancement. 140 Chang and Rochelle 141 investigated the impact of buffering by acetic acid and concluded that as little as 10 to 20mM of acetic acid would provide maximum mass transfer enhancement. Dibasic acid (commercially called AGS), which is a byproduct of adipic was also tested for buffering capabilities. It enhanced limestone scrubber performance as

35 effectively as adipic acid and is 40 -50 % cheaper than pure adipic acid. 142 Adipic acid, AGS, glutaric and succinic acid in concentrations of around 5 mM were found to increase the SO2 removal from 78% to 90%. 143 Limestone dissolution has also been shown to increase with the use of ammonium based salts. 144 145

3.2 Motivation in Estimating the Dissolution Rate of Mg(OH)2

As stated earlier, knowledge and predicting the dissolution rate of any sorbent is essential for the successful operation of any CO2/SO2 capture system. Very few dissolution rate studies have been performed for compounds other than limestone and other similar calcium based substances for Flue Gas Desulphurization (FGD) systems. At pH levels greater than 8, dissolution rate occurs slowly and via poorly defined mechanisms. FGD units operate at a pH ranging between 4 - 6 and therefore the dissolution rates at higher pH values has not received much study.

This study aims to investigate the impact of higher values of pH on dissolution rate of magnesium hydroxide. A system pH of 8.6 was used as a constant value in many experiments for estimation of dissolution rate. The reason behind this is that in the Mg(OH)2 – CO2 system, a pH of 8.6 has been estimated to be an optimum pH value, where the absorption of CO2 will take place with minimum solid (MgCO3) byproduct formation. The objective of this work was to study the impact of pH, temperature and stirring speed on magnesium hydroxide dissolution rate.

An experimental correlation was also developed which would predict the rate of dissolution under different conditions. The rate controlling mechanism was determined and conversion time equations were developed using the shrinking core model. Kinetic parameters like rate constant, order of reaction and activation energy was determined. The impact of particle size decrease on the dissolution rate was quantified using the shrinking core model.

3.3 Dissolution Rate Studies – Experimental Techniques

Dissolution rate studies can be divided into the following categories based on the experimental technique used

DISSOLUTION RATE STUDIES

FREE DRIFT Ph STAT

ROTATING ROTATING POWDERS POWDERS DISK/CYLINDER DISK/CYLINDER

Figure 9. Different techniques for estimating dissolution rate.

The free drift method has been used more extensively for geology related studies to determine the dissolution rates of calcite and aragonite, which are polymorphic forms of calcium carbonate. 130 146 It involves suspending the size fractioned particles in solution and measuring the pH/ calcium ion concentration change with respect to time. The pH stat technique involves maintaining the bulk solution at a constant pH value by the addition of a mineral acid. The amount of acid per unit time is used to calculate the dissolution rate. The main advantages of the pH stat technique over the free drift technique are: a) minimal variation in the state of solvent saturation 147, b) reaction is not interrupted during real time, c) other parameters such as temperature can be accurately controlled and thereby, we get more accurate values for the dissolution rate. Sometimes the solid is compressed into a solid disk and made to rotate in the

37 solution with/without pH control. This is done for those solids whose pH values are very difficult to control in the experimental range of operation. The rotating disk ensures that the mass transfer to and from the surface is well defined and can be altered by changing the rotating speed of the disk. 148 The disk surface can be pretreated and etched to produce reproducible surfaces. 149

The current study employs the pH stat technique for solids suspended in well stirred solutions.

This has been a technique commonly employed for estimating the dissolution rate of limestone for wet FGD applications and is more closely applicable to a practical scenario. The pH and temperature was controlled satisfactorily within the given range using pH stat/controller device and water bath respectively.

3.4 Materials and Methods

3.4.1 Experimental Setup

Magnesium hydroxide dissolution was estimated using a pH stat device [Cole Parmer

Chemcadet] where the dissolution rate was correlated to acid consumption required to maintain a preset value of pH. A pH stat device consists of three parts 1) pH Probe 2) pH meter 3) pH controller. The pH controller monitors the reading from the pH meter and when the set value is exceeded, it activates a pump (Walchem Diaphragm Pump, Cole Parmer) This pump delivers acid into the beaker, preventing further increase in pH. The weight of the acid being delivered is recorded by a balance. The pH reading and weight balance readings are monitored continuously by means of data acquisition software [Labview 2011 and Teraterm Pro]. The temperature of the beaker is kept constant by means of a hot water bath (0.5 C). The pH values were controlled within 0.2 during the entire experiment.

Acid Reservoir

Acid Feed Line Acid Pump Thermocouple

Scale pH Probe

Magnetic Stirring Bar

Data Acquisition System pH Stat Temperature Hot Plate/Stirrer Controller

Figure 10. Schematic of the dissolution rate experiment.

Pure distilled water was used in all the experiments. The pH probe was immersed inside a 700 ml beaker filled with 250 ml of distilled water. The pH, temperature and agitation speed were set to preset values and magnesium hydroxide powder (USP/FCC grade, Fisher Scientific) was added to the beaker. The slopes of the acid weight vs. time curves were calculated in the time interval of 5 to 20 minutes. This was done to ensure the pH values were stable and within the control range.. The surface area of magnesium hydroxide powders used was estimated using the

BET analysis method (TriStar 3000 V6.05).

3.4.2 Estimation of Solid Dissolution Flux

Magnesium hydroxide reacts with HCl in a 1:2 stoichiometric ratio as given below.

Mg( OH )2 2 HCl  MgCl 2  2 H 2 O (1)

The weight of acid vs. time curves was used to calculate the rate of acid consumption. The solid dissolution rate flux was calculated using the approach developed by Wang et al. 150

103 CS The pureacid mole flowrategmol minacid (/) (2) 

Because the solution pH is kept constant, all acid flow is neutralized by solid Mg(OH)2 dissolution, bringing in the d factor[Stoichiometric ratio of solid: acid] =1/2

103 C S d The solid Mg( OH )( dissolutionrategmol / )  secacid (3) 2 60

8 10 Cacid S d 2 Dissolutionrateof Mg( OH )(2 intermsof )( / flux Ngmol )s  c m sec (4) 6 BET Wsolid

Where Wsolid = Initial weight of the solid (g) and BET = surface area of the solid at the start

(m2/g), S = slope of the acid solution consumption (g/min),  = density of the HCl solution

3 (g/cm ) and Cacid = concentration of the HCl solution used for pH control (mol/liter)

3.5 Results and Discussion

Reagent grade magnesium hydroxide from Fisher Scientific was used for most of the experiments. The average BET surface area was estimated to be 5.07 m2/g. Hydrochloric acid was used as the medium to control pH because the product formed by reaction between

o Mg(OH)2 and HCl is MgCl2. MgCl2 is highly soluble in water (54 g/100 mL of water at 20 C) and does not interfere with the dissolution process of Mg(OH)2.

Typical pH and HCl consumption vs. time graphs are given below in figures 11 and 12

10.0

9.5

9.0

8.5

8.0

pH 7.5

7.0

6.5

6.0

5.5 0 5 10 15 20 25 30 Time (min)

Figure 11. Sample pH vs. time. (pH = 8.6, temperature = 52 oC and rpm = 700)

520

500

480

460

440

420

Weight (g) Weight Equation y = a + b*x 400 Adj R. Square 0.98976 Value Standard Error 380 Intercept 551.40173 0.40878 Slope -9.11336 0.0309 360 0 5 10 15 20 Time (min)

Figure 12. Weight of HCl solution vs. time. (pH = 8.6, temperature = 52 oC and rpm = 700)

Table 2. Summary of dissolution rate results

Temperature Slope[S] pH Rpm Solid Dissolution rate Grade of Mg(OH)2

(°C) (g/min) (gmol/cm2-sec)

22.5 4.07 7.6 700 4.621 * 10-11 Fisher

22.5 3.08 8.6 700 3.497 * 10-11 Fisher

22.5 1.26 9.6 700 1.430 * 10-11 Fisher

32.0 5.2 8.6 700 5.905 * 10-11 Fisher

42.0 7.14 8.6 700 8.108 * 10-11 Fisher

52.0 9.43 8.6 700 1.070 * 10-10 Fisher

22.5 0.77 8.6 700 1.008 * 10-11 Industrial

52.0 9.25 8.6 1110 1.050 * 10-11 Fisher

52.0 8.01 8.6 500 9.096 * 10-11 Fisher

22.5 3.74 8.6 1110 4.247 * 10-11 Fisher

3.5.1 Effect of pH

The effect of pH was investigated keeping the temperature at 23 °C, stirring speed at 700 rpm. acid and magnesium hydroxide concentration [0.01 M HCl and 0.01 M Mg(OH)2] constant.

Most dissolution experiments for limestone have been conducted in the pH range of 4 to 6.

Plummer et al. 125 demonstrated linear dependence on H+ ion concentration with the dissolution rate for limestone. The drop in dissolution rate with increasing pH is more drastic at lower pH values. Chan and Rochelle 151 showed that the dissolution rate of limestone at room temperature decreases by a factor of 6.7 from 4 to 5. It is evident from figure 13 that increasing the pH

42 results in a lower dissolution rate for magnesium hydroxide. Increasing the pH results in lowering the H+ ion concentration and consequently lower values of dissolution rates.

-23.8

-24.0

-24.2

-24.4 Model Polynomial

ln Ns Adj. R Square -- -24.6 Value Standard Error

Intercept -41.7748 -- -24.8 B1 4.7025 -- B2 -0.3075 --

-25.0

7.5 8.0 8.5 9.0 9.5 10.0 pH

Figure 13. ln Ns (solid dissolution flux) vs. pH.

3.5.2 Effect of Temperature

Figure 14 presents the effect of temperature on reagent grade Mg(OH)2, which was estimated keeping all other process variables constant. The Mg(OH)2 dissolution rate increases with increasing temperature and this effect seems to be quite significant. A diffusion controlled process is only moderately dependent on temperature and a chemical reaction controlled process is strongly dependent on temperature. The reason is that diffusion coefficient D is linearly dependent on temperature, while chemical reaction constant is exponentially related to temperature as given by the Arrhenius law. 152 Figure 15 shows that dissolution rates for

Mg(OH)2 at a pH of 8.6 demonstrates an excellent fit for the Arrhenius equation i.e. [Solid

-Ea/RT dissolution rate [Ns] = Ae ]. This result would be used in the experimental model for predicting the effect of temperature on the dissolution rate

1.1

) 1.0

-sec

2 0.9

0.8

0.7

gmole/cm

(

10 0.6

0.5

0.4

0.3

Ns (flux) x 10 Ns 20 25 30 35 40 45 50 55

o Temperature( C)

Figure 14. Solid dissolution flux vs. temperature.

-22.8 Equation y = a + b*x Adj. R Square 0.97039 -23.0 Value Standard Error Intercept -11.86073 1.16416 Slope -3592.88118 360.5278 -23.2

-23.4

-23.6

ln Ns

-23.8

-24.0

-24.2 3.0x10-3 3.2x10-3 3.4x10-3 3.6x10-3 -1 1/T (K )

Figure 15. Arrhenius plot of ln Ns (solid dissolution flux) vs. 1/T.

3.5.3 Effect of Agitation Speed

Figure 16 illustrates the effect of agitation speed on Mg(OH)2 under pH control and temperature of 52.oC The dissolution flux increases with increasing the stirring speed in the range from 500 to 700 rpm. Moo Young and Calderbank 153 had studied that impact of agitation on liquid mass transfer between liquid and suspended particles and arrived at the conclusion that convective mass transfer resulting from mixing is independent of particle size and is affected by specific agitation power. The increase in stirring speed increases the convective mass transfer between the solid Mg(OH)2 and liquid. The dissolution rate was found to be independent of the agitation speed after 700 rpm at higher temperatures. Agitation speeds below 500 rpm resulted in unstable pH control by the pH stat and therefore, most of the experiments were carried out at

700 rpm.

-22.94 Equation y = a + b*x -22.96 Adj R. Square 0.98833 Value Standard Error -22.98 Intercept -26.24628 0.24598 -23.00 Slope 0.50255 0.0385 -23.02 -23.04

ln Ns -23.06 -23.08 -23.10 -23.12 -23.14

6.1 6.2 6.3 6.4 6.5 6.6 ln 

Figure 16. ln Ns (solid dissolution flux) vs. ln ω.

3.5.4 Experimental Model

An experimental correlation was developed by Meserole et al. 154 for limestone dissolution

Ln Ns = a – b1 (1/T) + b2 (pH) + b3 ln (ω) + b4 ln (At) + b5 ln (Mg).

Where

Ns = the dissolution rate flux of calcium or magnesium

T = Solution temperature (K), pH = Solution pH and ω is the stirring rate in the reactor (rpm).

Based on the experimental data obtained above for magnesium hydroxide, a similar correlation was developed, which had the following form

3592.4 lnNspH 32.894  pH  0.3075( )2 4.7025( ) 0.5025ln ( ) T

This correlation is valid for the following ranges pH between 7.6 to 9.6

Temperatures ranging from 22 °C to 52 °C

Rpm ranging from 500 to 700

This dissolution flux model was compared with experimentally derived flux values and the results are in good agreement with each other as seen in figure 17.

)

-sec

gmol/cm

(

11 8

0 0 2 4 6 8 10 12

Experimental values of x 10 flux values Experimental 11 2 Predicted values of flux x 10 (gmol/cm -sec)

Figure 17. Comparison between the experimental and predicted values of the dissolution flux.

3.6 Kinetic Analysis and Shrinking Core Model

3.6.1 Selection of Rate Controlling Mechanism

The shrinking core model was chosen because of its ability to identify the rate controlling step and to provide essential kinetic parameters for the dissolution reaction like rate constant, order of reaction and activation energy. It is useful to describe situations in which solid particles are being consumed either by dissolution or reaction. The magnesium hydroxide particle is nonporous and spherical and dissolves according to shrinking core model. For this reaction, the following steps happen in succession: 155

a) Diffusion of the reactant H+ ions from the main body of the liquid through the liquid film

to the surface of the solid.

b) Reaction on the surface between H+ and the solid particle.

c) Diffusion of the reaction products from the surface back into the bulk of the liquid. The

particle shrinks continuously and there is no ash layer present.

The rate controlling step for dissolution reactions are surface chemical reaction controlled or liquid film diffusion controlled. The selection of the appropriate rate controlling mechanism is essential for the development of conversion equations using the shrinking core model. A common approach is to plot the kinetic data of rc/ro (radius of particle at time t/initial radius) vs. t/τ (time taken for particle to reach radius of rc/time taken for complete dissolution) and compare the results with predicted curves for reaction and diffusion control.

Magnesium hydroxide dissolution rates were tested without pH control at 52 oC at 700 rpm using with 15 and 30 strokes/min of 0.01 M HCl solution, till all the particles were completely dissolved. 700 rpm was chosen as an optimum value, since the dissolution rate does not increase further with an increase in mixing speed beyond 700 rpm. The two flow rates cover a range of the HCl solution used for all experiments under pH control. The shrinking particle size was measured at every 5 or 10 min over the entire dissolution process in the absence of pH control by using a laser particle size analyzer (Spectex PC-2000). Figure 18 and Figure 19 clearly show the predominance of chemical reaction control over film diffusion control.

1.0

0.8 Film diffusion control

0.6

c r 0.4 Chemical reaction control

0.2

0.0 0.0 0.2 0.4 0.6 0.8 1.0 t/

Figure 18. Rate controlling mechanism at temperature = 52 oC / 700 rpm / HCl stroke rate = 15.

1.0

0.8 Film diffusion control

0.6

c r 0.4

0.2 Chemical reaction control

0.0 0.0 0.2 0.4 0.6 0.8 1.0 t/

Figure 19. Rate controlling mechanism at temperature = 52 oC/ 700 rpm / HCl stroke rate = 30.

3.6.2 Conversion-Time Equations and Determination of Apparent Rate Constant

The conversion-time expressions for a chemical reaction controlled process are derived based on the principle that the rate is proportional to the available surface area of the particle. In the reaction Mg(OH)2 + 2HCl -> MgCl2 + 2 H2O, Mg(OH)2 is given the subscript ‘b’ and HCl as

‘a’.

11dNbdN rbk C  ba'' (5) bal 44r22 dtr dt where k" is the overall rate constant for the surface reaction (assumed to be first order for now)

+ 3 and Cal is the concentration of H ion in the liquid phase (mol/m )

155 The number of moles of Mg(OH)2 and HCl are related stoichiometrically as follows:

2 dNb   b dN a   b dV  4   b r c dr c (6)

Where b is the stoichiometric constant, ρb is the molar density of magnesium hydroxide

3 (mol/m of solid) and rc is the radius of the unreacted magnesium particle (m)

Substituting (6) in (5)

dr  c bk'' C (7) baldt

Integrating (7)

2 brr o c t '' (1 ) (8) k Cralo

The time required for complete dissolution is given at ro = 0

2 bor   '' (9) kCal

Fractional conversion and radius of the particle are related by

3 Volumeof unreacted core rc 1X B   3 (10) Total volumeof the particle ro

From (8) (9) and (10)

t (1  (1)  X ) 1/3 (11)  B

1/3 t krB(1  (1  X ) ) (12) kr is defined as the apparent reaction rate constant and is defined as

'' 1 kCal kr  (13) 2 bor

1/3 The values of kr can be calculated from the slope obtained by plotting time vs. (1-(1-XB) ) for various pH and temperature values. The slope is calculated from 5 to 20 minutes, which is the region where the pH is brought to a constant value.

Fractional conversion XB can be calculated using the formula

Weight of HCl consumed at timet X  (14) B Weight of HCl required for completedissolution

1/3 A sample plot of (1-(1-XB) ) vs. time, is shown in Figure 20. It shows a good fit with a

2 coefficient of determination (R ) value of 0.989. The value of kr for the absorber operating at a

o -1 -1 pH of 8.6 and temperature of 52 C was estimated to be 1.1×10 min . The effect of temperature on dissolution rate can provide information on the dominant reaction mechanism for the process. Activation energy values at the desired operating pH of 8.6 can be obtained by

-Ea/RT using the Arrhenius equation kr = ko e , where ko is the pre-exponential factor, Ea is the activation energy (J/gmol), R is the gas constant (8.314 J/gmol K) and T is the temperature (K).

When the temperature was raised from 22 C to 52 C, the apparent reaction rate constant (kr) increased by a factor of 5.4 as shown in the Arrhenius plot in Figure 21. Diffusion controlled

51 reactions have activation energies lower than 20 kJ/gmol and surface controlled reactions have

137 156 activation energies between 40 to 80 kJ/gmol. The activation energy for Mg(OH)2 dissolution at a pH of 8.6 was calculated to be 42 6 kJ/gmol. This was comparable with the values obtained (45 kJ/gmol) at similar temperatures for a South African magnesium-based material, whose rate determining step was found to be controlled by surface chemical reaction.

157 This result proves that surface reaction plays a dominant role in the dissolution of Mg(OH)2 at the high pH regime.

0.25 R2 value = 0.989 slope value (5 to 20 min) = 0.01081 0.20

) 0.15

1/3

)

1-X

( 0.10

(

0.05

0.00 0 5 10 15 20 25 Time (min)

1/3 o Figure 20. Sample plot of (1-(1-XB) ) vs. Time (min) for pH = 8.6 and temperature = 52 C.

-4.0 -4.4 52 oC

-4.8 42 oC

-5.2 32 oC

) -5.6

-1 -6.0 o min Equation y = a + b*x 22 C

( r -6.4 Adj. R-Square 0.93444 Value Standard Error

ln k -6.8 -- Intercept 11.04694 2.44303 -7.2 -- Slope -5042.64954 762.29901 -7.6 -8.0 0.0030 0.0031 0.0032 0.0033 0.0034 0.0035 -1 1/T(K )

o Figure 21. Activation energy calculation using kr at pH = 8.6 and temperature = 52 C.

3.6.3 Order of Reaction and Overall Activation Energy for Mg(OH)2 Dissolution

+ The values of kr calculated from experimental data at same temperature and different H ion concentrations should yield the same value of k", which is the overall reaction constant. It is

" found however, that the value of k differs with changes in the Cal value. This implies that rate of dissolution is not directly proportional to the concentration of H+ ion in the liquid and the order of reaction is not first order.

Equation 9 was modified as follows:

'' n kCal kr  (15) 2 bor

Where n is the fractional order of reaction. Taking natural logarithm on both sides of equation (15)

k '' lnkr ln z  n ln C al where z  (16) 2bor

-3.9 Unit for k '' = (min-1 (mol/m2) ((mol/m3) -n) s l o -4.2 52 C n = 0.31 -4.5 k'' = 1.427 *10-10 42 oC -4.8 o n = 0.23 32 C -5.1 -11 ) k'' = 3.420 *10

-1 -5.4

min n = 0.20 ( -5.7 r o k'' = 1.658 *10-11 22 C -6.0

ln k -6.3 n = 0.22 k'' = 8.256 * 10-12 -6.6 -6.9 -7.2 -16 -15 -14 -13 -12 -11 -10 ln C (mol/m3) Al

Figure 22. Plot of ln kr vs ln Cal to determine n and k” at different temperatures.

o The value of n was calculated by plotting the kr values at four different temperatures (22 C, 32 o o o + C, 42 C and 52 C) and different H ion concentrations calculated from the constant pH values

(figure 22). The reaction order (n) values determined were 0.22 at 22 C, 0.20 at 32 C, 0.23 at

42 C, and 0.31 at 52 C, from which the arithmetic average n value was found to be 0.24. The

-1 2 3 -n reaction rate constant (k with a unit of min (mol/m )s(mol/m )l ) determined from the intercept values were 8.26×10-12 at 22 C, 1.66×10-11 at 32 C, 3.42×10-11 at 42 C and

1.43×10-10 at 52 C. From the intrinsic/overall reaction rate constant, the true activation energy

-1 2 3 -n and frequency factor were found to be 7611 kJ/gmol and 220 min (mol/m )s(mol/m )l from the Arrhenius equation (figure 23). These results are in agreement with values reported for other calcium and magnesium based minerals.

-22

)

-n

l Equation y = a + b*x ) o ) Adj. R-Square 0.93912

3 52 C Value Standard Error -23 -- Intercept 5.39566 4.3037 -- Slope -9192.1073 1336.83949

mol/m

(

) 2 -24 42 oC

mol/m

(

-1 o -25 32 C

(min) ( 22 oC -26 ln k'' 0.0031 0.0032 0.0033 0.0034

-1 1/T (K )

Figure 23. Plot of ln k’’ vs. 1/T to determine the overall activation energy.

Fredd and Fogler studied the dissolution rates of calcite in acetic acid solutions and concluded that above 3.7, the dissolution rate is influenced primarily by chemical reaction. The dissolution of mineral oxides in an aqueous environment has been studied for its potential applications in environmental remediation, corrosion, and drug design. Previous studies have shown that the magnesium oxide dissolution exhibits a fractional order n which is frequently in the range between (0

(ranging from 2 - 4) and different acid anions also showed the predominance of surface reaction control and that a fractional order for the reaction with H+ ions coupled with high values of

55 activation energies serve as an indicator for this type of reaction control. 161 The dissolution kinetics of chrysotile (magnesium silica based hydroxide (white asbestos) – Mg3Si2O5(OH)4) was investigated due to public health significance of contamination in drinking water and exhibited a fractional dependence on pH (n = 0.24) at pH values ranging from 7 to 10 at

25 C. 162 Fruhwith et al. 163 reported an activation energy value of approximately 70 kJ/gmol for MgO dissolution at pH values >7 with temperatures ranging from 25 to 75 C. However, there was no experimental justification provided about the impact of pH on the dissolution rate and the information was provided about the nature of the reaction control regime was not clear.

Nevertheless, these findings substantiate the existence of a fractional order, high activation energies and dominance of surface chemical reaction for the dissolution of magnesium hydroxide at higher pH regimes, which has not been explored much in previous studies.

3.6.4 Modified Rate of Dissolution Expression Based on the Shrinking Core Model

The fractional order of reaction results in modification of equation (5), as follows

11dN bdN rba    bk'' C n (17) b44r22 dt r dt al

1 dNb '' n rbkbal C  (18) Sa dt

n n n Cbl bC al  C al  b C bl (19)

The flux expression for Mg(OH)2 dissolution in terms of concentration of the solid is given by

1dN E  9,192   gmol  rba  b1n k '' C n  b 1 n kexp C n  130exp  C 0.24 (20) b bl0  al bl 2 Sa dt RT T  m min 

3.6.5 Effect of Decreasing Particle Size During the Dissolution Reaction.

From the shrinking core model we have

n k'' CAl t rrco(1) (21) 2 bor

The radius of the particle at any time t can be calculated by the above formula and using the values of n and k" calculated previously. It compares well with the measured data calculated by equation 22 for the intrinsic dissolution period (20 minutes) as seen in figure 24.

rc 1/3 (1) X B (22) ro

An experiment was run for 60 minutes at pH = 8.6 and temperature = 52 οC. The values obtained were verified experimentally by taking samples out at fixed time from the dissolution experiment. The samples were washed with water and the particle size was analyzed using the particle size analyzer with a base dilution of 107. Figure 25 shows a representative particle size distribution for the sample taken at 0 min. The system deviates from the intrinsic region after 20 minutes as seen by the figure 26. This is due to the system attaining an equilibrium value beyond which the Mg(OH)2 does not dissolve further.

1.0 Measured data Shrinking core model

0.9

0.8

r 0.7

0.6

0.5 0 4 8 12 16 20 Time (min)

Figure 24. Comparison of measured data with shrinking core model at pH = 8.6 and temperature = 52 oC.

200 (Average particle size - 5.71m)

150

100

Counts 50

1 - 2 2 - 4 4 - 8 8 - 16 16 - 31 31 - 63 63- 128 Particle size distribution (m)

Figure 25. Particle size distribution for sample taken at 0 min.

6.0 Measured data 5.5 Particle size analyzer values

) Theoretical SCM values

m 5.0



( 4.5

4.0

3.5

3.0

2.5

Average particle size size particle Average

2.0 0 10 20 30 40 50 60 Time (min)

Figure 26. Average particle size vs. time using SCM model and experimental data.

The same analogy for particle radius can be extended to the area occupied by a single particle. The dissolution rate decreases with decreasing particle size and for a single particle, the area at any time t can be calculated by

Ar cc (23) Aroo

Substituting (23) in (22)

2/3 AAXc o(1 B ) (24)

It can be seen from figure 27 that the about 60% of the area of the particle is unreacted at time t =

20 min and the decrease in area is linear. But as time progresses the decrease in the area of the particle affects the dissolution rate.

1.1

1.0

0.9

0.8

0.7

0.6

0.5

0.4

0.3

0.2 0 5 10 15 20 25 30 35 40 45 50 55 60 65

Fraction of unreacted area of a particle of a area of unreacted Fraction Time (min)

Figure 27. Fraction of unreacted area vs. time.

The slope of HCl consumed vs. time deviated increasingly away from the linear region with time as shown by figure 28. The value calculated upto 20 minutes is called the intrinsic dissolution rate or the maximum dissolution rate, upto which the dissolution rate is unaffected by the shrinking particle size. The duration of this period will depend upon operating conditions like pH, temperature, concentration and stirring/agitation speed. The change in the average size of the particle with respect to normalized time is illustrated by figure 29. It can be seen that the when the average particle size gets reduced to 60% of its initial value, the system deviates from its intrinsic chemical reaction control regime, which would result in a decrease in the dissolution rate. The rate of CO2 removal would be maximum if the Mg(OH)2 slurry in the CO2 absorber is made to dissolve at this intrinsic rate. Particle size analysis of Mg(OH)2 particles intermittently would be a good tool to predict whether the slurry needs to be recirculated or replaced for optimum performance.

600

550

500

450

400

350

300

HCl consumpition (g) HCl consumpition 250

200 0 10 20 30 40 50 60 Time (min)

Figure 28. Weight of HCl consumed vs. time for pH = 8.6, temperature = 52 οC for 60 min.

1.0

0.8 Film diffusion control

0.6

c r 0.4 Chemical reaction control 0.2

0.0 0.0 0.2 0.4 0.6 0.8 1.0 t/

Figure 29. Plot showing the effect of particle size on the rate controlling mechanism

3.6 Conclusion

The dissolution rate of magnesium hydroxide under different operating conditions of pH, temperature and stirring rate was measured using a pH stat apparatus. An experimental correlation was developed which can predict intrinsic Mg(OH)2 dissolution rates as a function of pH, temperature and stirring rate. The dissolution of magnesium hydroxide particles was modeled using the shrinking core model and the rate of dissolution was found to be dependent on surface chemical reaction control. The value of the apparent reaction constant was determined to be 1.1×10-1 min-1 for a pH value of 8.6 and temperature of 52 oC, which are the absorber conditions in the current system. The effect of temperature on the apparent rate constants was found to follow Arrhenius equation and was estimated to be 42 6 kJ/gmol at a pH value of 8.6.

+ The dissolution of Mg(OH)2 was found to exhibit fractional dependence on the H ion concentration. The order of reaction was found to have a range of 0.20

o for Mg(OH)2 under the absorber conditions of pH = 8.6 and temperature = 52 C would lead to deviation from the intrinsic dissolution rate, controlled by surface chemical reaction control.

CHAPTER 4. KINETICS OF MgCO3 FORMATION.

4.1 Introduction

Magnesium carbonate is a crystalline solid formed when magnesium hydroxide reacts with CO2. Figures 30 and 31 show the equilibrium distribution of carbonate and magnesium species which are distributed as a function of pH.

100

Legend Mg(2+) MgOH+ 80

m MgHCO3+

s MgCO3(aq)

M 60

t 40

P 20

0 2 4 6 8 10 12 14 pH Figure 30. Distribution of dissolved magnesium species as a function of pH.

100

C Legend

60 -2 d CO3

l H2CO3 (aq)

s - i HCO3

l + MgHCO3 a t 40

c r

P 20

0 2 4 6 8 10 12 14 pH

Figure 31. Distribution of dissolved carbonate species as a function of pH.

This was estimated by running MINEQL which is a chemical equilibrium modeling package for phase equilibrium determination. At a pH between 8 and 9 which is the optimum pH range, the ionic species are dominated by the bicarbonate ion. This would ensure that the formation of magnesium carbonate is minimized. Magnesium carbonate formation would reduce the lifetime for solvent in the chemical cycle between the absorber and the stripper. Dissociation of MgCO3 would require very high temperatures [>250 oC], as shown in the following TGA diagram for pure MgCO3.

100

Weight % Weight 50

30 0 100 200 300 400 500 600 700 o Temperature ( C)

Figure 32. TGA curve for MgCO3.

Previous studies have been done in estimating the crystal growth rate of calcium sulfite and calcium sulfite hemihydrate, which are the main reaction products of a flue gas desulphurization unit. Calcium sulfite crystallization affects several factors such SO2 absorption, limestone dissolution and solution composition, in a manner analogous to MgCO3. Tseng and

Rochelle et al. 164 studied the kinetics of calcium sulfite hemihydrate growth using a pH stat

64 apparatus in an aqueous solution of pH 3.5 to 6.5 and found that the growth was a function of relative supersaturation and strongly inhibited by the dissolved sulfate. Boke et al. 165 developed a quantitative analysis using FTIR to distinguish between limestone(CaCO3), calcium sulfite hemihydrate (CaSO3. 0.5 H2O) and gypsum(CaSO4. 2H2O). The method employed in this study involves using TGA-MS and HCl titration to estimate the distribution of carbon among the solid and liquid phases. The effects of pH control and the heating, on the kinetics of magnesium carbonate formation were investigated. Figure 30 shows that maintaining a pH of 8.6 is conducive to prevent the formation of magnesium carbonate and ensuring that the carbon stays in the liquid phase and interacts with the magnesium in the liquid phase. This would help in the regeneration of the magnesium hydroxide solution in the stripper phase. Also the effect of higher temperatures employed [52 C and 65 C] in the process needed to be investigated.

4.2 Experimental Setup and Procedure

A 1M Mg(OH)2 was prepared by adding 29.15 g of Mg(OH)2 to 500 ml of water. A total inorganic concentration of X tic = 0.01 was simulated by adding 0.56 moles of sodium bicarbonate to the Mg(OH)2 solution. Sodium bicarbonate [NaHCO3] was used as a source of carbonate and bicarbonate ions to simulate CO2 in the liquid phase. A pH stat with a set point of

8.6 was connected to pump containing HCl of 1N. The pH stat monitors the pH in the batch setup and prevents it from going above 8.6. A temperature probe was used to monitor the temperature of the solution and a hot plate/stirrer was used to set the temperature of the batch system.

Figure 33. Experimental schematic for kinetic analysis.

Samples were taken out at different time intervals and immediately centrifuged for the shortest possible time to separate the solid from the liquid. The liquid was filtered with a 0.45 μm filter and analyzed for liquid carbon concentration by HCl titration method. The solids were dried under vacuum at 60 for 24 hrs. till the solid became dry. The solids were then analyzed for carbon content using the TGA–MS.

4.2.1 Solid Analysis using the Thermo-gravimetric Analyzer - Mass Spectrometer.

The solid samples were analyzed for carbon content by using a Thermo gravimetric analyzer [TA Instruments TGA Q5000IR] coupled with a Mass Spectrometer [Pfeiffer-Vacuum

Thermostar]. Percent weight loss and detected spectra vs. temperature/time were used to quantify the CO2 present in the solid sample. Oxygen was used as a carrier gas with the high flow rate of 100 ml/min to ensure complete combustion with approximately 10 ~ 25 mg of sample. Ramp rate used was 5 ˚C/min upto a target temperature of 800 ˚C. The calibration curve was obtained using known samples of calcium oxalate powder. The calibration curve for

CO2 and the TGA-MS curves for a sample experiment are shown below in figures 34, 35 and 36 respectively.

Equation y = a + b*x

Ad. R Square 0.99033

Weight (mg) Weight

2 Value Standard Error

Intercept 0.3115 0.23086 Slope 4.05985 0.20035

0 0.0 0.5 1.0 1.5 2.0 Area (nA.min)

Figure 34. Calibration curve for CO2 measurement using the TGA-MS.

Figure 35. Sample TGA-MS curve for pH control and temperature 65 oC at 0 min.

Figure 36. Sample TGA-MS curve for pH control and temperature 65 oC at 15 min.

4.3 Results and Discussion

The mass balance tables with pH control using HCl and without pH control at three different temperatures are presented below. Bar graphs show the partitioning of the carbon

- 2- between the bicarbonate ion in the liquid [HCO3 ], carbonate ion in the liquid [CO3 ] and the solid carbon over different time intervals.

4.3.1 No pH control and Temperature of 52 C

Table 3. Mass balance for no pH control and temperature of 52 oC

Sample Time(min) Carbon in Carbon in Solid phase Actual Mass 2- - No CO3 (g) HCO3 (g) Carbon(g) Balance(g)

1 1 0.16 2.89 0.23 3.28

2 3 0.22 2.84 0.50 3.56

3 5 0.32 2.48 0.88 3.68

4 7 0.38 2.32 0.84 3.54

5 10 0.44 2.20 0.80 3.44

6 15 0.48 2.19 0.75 3.42

Theoretical – 3.36 g carbon input, Error range – 2% to 8%

solid carbon (g) HCO- carbon(g) 3 CO2- carbon(g) 3

3.5

3.0

2.5

2.0

1.5

1.0

Carbon Content (g) Content Carbon 0.5

0.0 1 3 5 7 10 15 Time (min)

Figure 37. Carbon content vs. time for no pH control at 52 °C.

The mass balance table (table 3) and the bar graphs (figure 37) show the increasing evolution of carbon in the solid phase with time. The drop in the bicarbonate concentration from

1 to 15 minutes is around 20%. It is also seen that the carbonate concentration in the liquid phase is increasing with time.

9.4 9.3 9.2 9.1 9.0 8.9

pH 8.8 8.7 8.6 8.5 8.4 0 2 4 6 8 10 12 14 16 Time (min)

Figure 38. pH vs. time graph for no pH control at 52 °C.

There was a decrease in the initial pH of the Mg(OH)2 from around 10 to 9.2, due the increase in temperature from 24 °C to 52 °C (figure 38). The pH dropped quickly at the start, because sodium bicarbonate, being soluble in water, released H+ ions and brought down the pH.

+ Then the pH increased slowly till around 3.5 minutes. In this region, the H ions reacted with the

- OH ions (released from the dissociation of Mg(OH)2) and formed water. This region also included the initiation of solid carbon formation, which continued into the next zone upto 10 minutes. After this period, the pH increased rapidly and went back up to the starting pH. This

70 was mainly due to the dissolution of magnesium hydroxide, which released OH- ions, thereby increasing the basicity of the solution.

4.3.2 pH control at 8.6 and temperature of 52 C

Table 4. Mass balance for pH control at 8.6 and temperature at 52 oC

Sample Time(min) Carbon in Carbon in Solid phase Actual Mass 2- - No CO3 (g) HCO3 (g) Carbon(g) Balance(g)

1 1 0.13 3.04 0.33 3.50

2 3 0.19 2.82 0.66 3.67

3 5 0.22 2.40 0.94 3.56

4 7 0.17 1.90 1.01 3.08

5 10 0.14 1.56 1.15 2.85

solid carbon (g) HCO- carbon(g) 3 CO2- carbon(g) 4.0 3

3.5

3.0

2.5

2.0

1.5

1.0

Carbon Content(g) Carbon 0.5

0.0 1 3 5 7 10 Time (min)

Figure 39. Carbon content vs. time for pH control at 52 °C.

The data for the pH control and temperature 52 C (table 4 and figure 39) showed a dramatic decrease in the bicarbonate carbon concentration, when compared to the no pH control case. The drop in the bicarbonate carbon concentration from 1 minute to 10 minutes is around

48% for the pH control case, compared to 23% bicarbonate carbon concentration drop, for the same time period, in the experiment with no pH control. This drop in bicarbonate concentration is explained by the following sequence of steps.

When pH control was applied, most of the carbon remained in the liquid phase as bicarbonate ion, and reacted with magnesium ions to form magnesium bicarbonate

[Mg(HCO3)2]. In addition, the acid used for pH control was of high concentration [1N] and this might have caused more extensive dissolution of the magnesium hydroxide, leading to greater production of magnesium ions and higher reactivity with the bicarbonate ions in the solution. 166

The drop in bicarbonate concentration evident during liquid phase titration occurs due to two possible reasons: a) The [Mg(HCO3)2] in the liquid phase reacts with HCl giving rise to soluble

MgCl2 and CO2 gas, b) Some portion of the magnesium bicarbonate also ends up with the solids when the supernatant liquid was separated for liquid phase titration. It was reported that magnesium bicarbonate can be converted into a hydrated form of magnesium carbonate, when the slurry was separated and dried under vacuum. 167 This was further proved by the appearance of more carbon in the solids upon analysis. However, this experiment proves that under

o conditions of pH control at 52 C, the formation of [Mg(HCO3)2].is enhanced ,which appears as a drop in bicarbonate concentration during the mass balance analysis.

10.0

9.8

9.6

9.4

9.2

pH 9.0

8.8

8.6

8.4 0 2 4 6 8 10 12 14 16 Time(min)

Figure 40. pH vs. time for pH control at 52 °C.

The pH stat maintained the pH at 8.6 throughout the experiment.

4.3.3 No pH control and temperature of 65 C

Table 5. Mass balance for no pH control and temperature of 65 oC

Sample Time(min) Carbon in Carbon in Solid phase Actual Mass 2- - No CO3 (g) HCO3 (g) Carbon(g) Balance(g)

1 1 0.15 3.28 0.44 3.87

2 3 0.28 2.88 0.90 4.06

3 5 0.36 2.66 1.10 4.12

4 7 0.42 2.52 1.17 4.11

5 10 0.46 2.48 1.10 4.04

6 15 0.48 2.48 1.15 4.11

solid carbon (g) HCO- carbon(g) 3 CO2- carbon(g) 4.5 3

4.0

3.5

3.0

2.5

2.0

1.5

1.0

Carbon Content (g) Content Carbon 0.5

0.0 1 3 5 7 10 15 Time(min)

Figure 41. Carbon content vs. time for no pH control at 65 °C.

The increase in temperature resulted in the accelerating the kinetics of solid carbon formation as seen in figure 41 and table 5. When compared to the sample taken at 3 minutes for the experiment performed with no pH control and 52 °C, there was almost a 50% increase in the solid carbon formation for no pH control and 65 °C. This effect reduced during the entire experimental run and the percentage difference in the solid carbon formation came down to 34%, when the samples taken at 15 minutes were compared.

9.4

9.2

9.0

8.8

8.6

0 2 4 6 8 10 12 14 16 Time (min)

Figure 42. pH vs. time for no pH control at 65 °C.

The pH of the solution decreased further from its equilibrium room temperature value when the temperature is increased to 65 °C as seen in figure 42. The time period where the pH increased slowly was much shorter for this experiment when compared to the experiment conducted at 52 °C. This occurred because with the increased temperature, the reaction between the various ions occurred almost instantaneously.

4.3.4 pH control at 8.6 and temperature of 65 C

Table 6. Mass balance for pH control at 8.6 and temperature of 65 oC

Sample Time(min) Carbon in Carbon in Solid phase Actual Mass 2- - No CO3 (g) HCO3 (g) Carbon(g) Balance(g)

1 1 0.12 3.14 0.63 3.89

2 3 0.14 2.60 1.10 3.84

3 5 0.17 2.13 1.68 3.98

4 7 0.12 1.98 1.97 4.07

5 10 0.12 1.62 1.92 3.64

6 15 0.11 0.90 2.66 3.67

solid carbon (g) HCO- carbon(g) 3 CO2- carbon(g) 4.5 3

4.0

3.5

3.0

2.5

2.0

1.5

1.0

Carbon Content (g) Content Carbon 0.5

0.0 1 3 5 7 10 15 Time (min)

Figure 43. Carbon content vs. time for pH control at 65 °C.

Same trend was observed as in the previous experiment with 52 °C, with and without pH control as seen in table 6 and figure 43. The effect was even more pronounced at higher temperatures as shown in these experiments with 65 C. After 15 minutes, the percentage of

76 available bicarbonate ion was almost 75% more in the no pH control case. This validates our hypothesis that during the process of pH control, the bicarbonate ion predominates and leads to an increased formation of Mg(HCO3)2. This increase is also helped by the enhanced dissolution of Mg(OH)2 on reacting with HCl. The higher values of magnesium bicarbonate formed during the course of the reaction appears during the liquid phase carbon analysis in the form of bicarbonate concentration and these values are lesser than the value obtained at the same time for the non pH control case. This drop can be attributed to the conversion of Mg(HCO3)2 to CO2 during the course of the reaction with HCl or to hydrated magnesium carbonate during the solid phase carbon analysis, which is substantiated by the increase in solid carbon content.

10.0

9.8

9.6

9.4

9.2

pH 9.0

8.8

8.6

8.4 0 2 4 6 8 10 12 14 16 18 Time(min)

Figure 44. pH vs. time for pH control at 65 °C.

The pH stat kept the pH stable at 8.6 without any fluctuation.

4.3.5 No pH control and room temperature [24 °C]

Table 7. Mass balance for no pH control and temperature of 24 oC

Sample Time(min) Carbon in Carbon in Solid phase Actual Mass 2- - No CO3 (g) HCO3 (g) Carbon(g) Balance(g)

1 1 0.24 3.16 0.35 3.75

2 3 0.26 3.10 0.32 3.68

3 5 0.26 3.10 0.44 3.80

4 10 0.27 2.90 0.39 3.56

5 20 0.36 2.50 0.92 3.78

6 30 0.46 2.28 1.20 3.94

solid carbon (g) HCO- carbon(g) 3 CO2- carbon(g) 4.0 3

3.5

3.0

2.5

2.0

1.5

1.0

Carbon Content(g) Carbon 0.5

0.0 1 3 5 10 20 30 Time(min)

Figure 45. Carbon content vs. time for no pH control at 24°C.

The rate of solid formation (table 7 and figure 45), when compared to the experiments run at elevated temperatures and no pH control was slow. For example, the amount of solid

78 carbon formed at 10 minutes is 64% less than solid carbon formed at 65 °C for the same time period.

10.2

10.0

9.8

9.6

9.4

pH 9.2

9.0

8.8

8.6

0 5 10 15 20 25 30 35 Time(min)

Figure 46. pH vs. time for no pH control at 24 °C.

The initial pH of Mg(OH)2 slurry solution is ~10, and starts to drop due to the dissolution

 2 and equilibration of HCO3 in NaHCO3 with CO3 as shown in figure 46. After some of is consumed to form MgCO3 (which is warranted by a constant pH region between 5 and 20 minutes), OH- starts to be released to return to the initial pH value, ~10.

4.3.7 pH control at 8.6 and room temperature [24 °C]

Table 8. Mass balance for pH control and temperature of 24 oC

Sample Time(min) Carbon in Carbon in Solid phase Actual Mass 2- - No CO3 (g) HCO3 (g) Carbon(g) Balance(g)

1 1 0.14 2.88 0.24 3.26

2 3 0.12 2.86 0.26 3.24

3 5 0.10 2.86 0.28 3.24

4 10 0.15 2.77 0.31 3.23

5 20 0.11 2.17 0.80 3.08

solid carbon (g) HCO- carbon(g) 3 CO2- carbon(g) 3.5 3

3.0

2.5

2.0

1.5

1.0

Carbon Content (g) Content Carbon 0.5

0.0 1 3 5 10 20 Time(min)

Figure 47. Carbon content vs. time for pH control at 24 °C.

The data for the no pH control and room temperature showed a gradual increase in the solid carbon concentration with time and after 30 minutes almost 30% of the carbon was in the

80 solid phase. But when pH was controlled to 8.6 (table 8 and figure 47), the amount of bicarbonate carbon decreases, compared to the no pH control case. This decrease is not substantial, when compared to the other experiments at higher temperatures. Consequently, the difference in the amount of carbon in the solid is not significant, despite pH control. This is because the process takes place at room temperature and thereby the net amount of solid formed is also less, when compared to the experiments at higher temperature.

10.2 10.0 9.8 9.6 9.4 9.2

pH 9.0 8.8 8.6 8.4

0 2 4 6 8 10 12 14 16 18 20 22 Time(min)

Figure 48. pH vs. time for pH control at 24 °C.

4.4 Conclusion

A mass balance analysis was carried out to estimate the distribution of carbon among the different phases for the Mg(OH)2-H2O-NaHCO3 system and the mass balance was closed satisfactorily. The effect of temperature on MgCO3 formation was investigated and the results showed that the increase in reaction temperature would lead to an increase in MgCO3 formation.

MINEQL software predicted that for minimizing the carbonate formation, the system needs to be

81 operated at a pH of 8.6. The impact of pH control at three different process temperatures was investigated and it was found that pH control at 8.6 leads to the formation of magnesium bicarbonate, which is highly beneficial for the regeneration of the Mg(OH)2 solvent. The drop in bicarbonate concentration in the liquid over time and the concurrent increase in solid carbon justifies our results. But as seen even with pH control, over the course of time, the system tends to move towards the more stable state of solid carbonate formation. This can be offset by utilizing chemicals that might block that growth of carbonate crystal, employing shorter contact times in the absorber or by regular replenishment of the solvent.

CHAPTER 5. VAPOR-LIQUID-SOLID EQUILIBRIUM (VLSE) STUDIES

5.1 Introduction

Vapor-liquid-solid equilibrium (VLSE) data for the CO2–water–Mg(OH)2 system at various temperatures, pressures and concentrations have a vital role in the design and

168 optimization of the flue gas treating process[dash paper]. The equilibrium solubility of CO2 determines the amount of solvent feed solution that needs to be circulated to treat a given feed

169 gas. It also determines the maximum concentration of CO2 which can be left in the regenerated solution in order to meet the treated gas specifications. 170 Physical and chemical equilibria also play a crucial role in defining the boundary conditions for partial differential equations that describe the mass transfer coupled with chemical reactions. VLE data would also be essential to estimate solvent degradation at elevated temperatures. There are hardly any studies on VLSE which were found in literature which exactly mirror the chemistry that is found in the current system of interest i.e. [CO2–water–Mg(OH)2]. Most studies have been performed on amines like MEA (Monoethanolamine), DEA (Diethanolamine) and MDEA

(Monodiethanolamine) piperazine and mixtures of amines with piperazine. This data has been summarized in the form of a table (Table 9). This table contains information about the system under consideration, reaction conditions and additional information investigated by various authors.

Table 9. Literature review of VLE estimation

Source System under Temperature Pressure Additional information Consideration range range (partial pressure of

CO2 )

Ma’mun 30 % mass MEA and 50 MEA – [120 MEA - 7 to Enthalpy of solution of 171 o et. al % mass MDEA C] 192 kPa CO2 in the aqueous MDEA – [55, MDEA – 66 MDEA solution was 70 and 85 oC] to 813 kPa estimated

Aronu et 15, 30, 45 and 60 mass % 40 – 120 oC MEA – al. 172 MEA 0.0035 kPa to 1243.41

Chakma MDEA [4.28 M and 1.69 100 oC to 200 MDEA – et al. 173 M] oC 138 to 4930 kPa

Liu et Piperazine (0.17 kmol/m3 30 oC to 90 oC 13.16 to al. 174 to 1.55 kmol/m3) + 935.3 kPa MDEA (1.35 to 4.77 kmol/m3)

Derks et Piperazine [0.2 to 0.6 M ] 25,40 and 70 0.3 to 111 al. 175 oC kPa

Dash et Piperazine [0.2 to 4.5 M ] 25, 35, 45 and 0.1 to 1500 al. 168 55 oC kPa

Shen et 12 wt % MEA + 18 % 40, 60,80, 100 1 to 2000 al. 176 MDEA, 24 wt% MEA + oC kPa 6 wt% MDEA

Austgen 2.5 kmol/m3 MEA, 2.0 40 oC and 80 0.05 to et al. 177 kmol/m3 & 4.28 kmol/m3 oC 309.3 kPa MDEA , 2.0 kmol/m3 MDEA + 2.0 kmol/m3 DEA

Puxty et 1 mol/L of MEA, AMPD, 40 oC 0.3 to 908 Large Scale screening al. 178 two novel amines (A1 and kPa process was carried out A2) in identifying amines having best CO2 absorption.

Dash et 2-amino-2 –methyl-1- 25 oC to 55 oC 0.4 to 1500 al. 179 propanol (2.2 to 4.9) mol kPa dm-3

The parameters in vapor-liquid equilibrium measurement are temperature, pressure, xi[liquid phase mole fraction of component i] and yi[vapor phase mole fraction of component].

The value of i ranges from 1 to n where n is the number of components. The most commonly fixed variable in VLE measurement is temperature, which is held constant by means of a water bath or a temperature controller. Additionally fixed variables may include the amount of solvent, amount of component loading in the liquid phase and partial pressure of volatile components in

85 the vapor phase. The most widely employed method for studying VLE is the static method, in which a thermostated system having a set empirical composition, is brought into thermodynamic equilibrium and the phase compositions and the pressure are directly determined. 180 This usually achieved by bringing an aqueous amine solution(solvent) in contact with CO2 or H2S in a cell and waiting till equilibrium is achieved. The indication for onset of equilibrium can be a constant value pressure or pH. Some experimental designs also allow for analysis of the liquid and vapor phases of the desired component by sampling. The current study uses a static method for measurement of vapor-liquid-solid equilibrium data.

5.2 Experimental Setup and Procedure

A high pressure Parr reactor [model 4768] with a working volume of 1 liter was used.

This has an inlet valve with a dip tube for the gas which can double up as a liquid collection valve and an outlet valve. The relief valve was preset to 200 psi. A lecture bottle with a volume of 440 ml was used to transfer the CO2 to the reactor. A magnetic stirrer was used to keep the slurry in a well-mixed state. A thermocouple connected to a temperature controller inserted into the reactor to monitor the temperature. Heating plate/water bath were used to raise the temperature of the reactor when needed.

Pressure Inlet valve for gas gauge P Gas inlet line (0-200)psi

Gas bag

thermocouple

Pressure P gauge

Water Bath

High pressure reaction vessel T Temperature Lecture controller bottle Magnetic stirrer Figure 49. Schematic diagram of the experimental setup.

5.2.1 Experimental Procedure

The lecture bottle was a small pressurized gas cylinder, used to transfer CO2 gas from a main CO2 gas cylinder. It was equipped with a pressure gauge with a range of 0 to 200 psig.

The weight of the lecture bottle was recorded before transferring CO2 gas to the Parr reactor. A magnesium hydroxide slurry solution was added to the vessel. The vessel was closed and connected to a vacuum pump in order to vacuum residual air and water vapor occupied in the head space of the vessel. After a vacuum was created, CO2 gas was transferred to the reactor containing Mg(OH)2 slurry solution until the pressure reached a desired total pressure. As CO2 gas dissolved in the Mg(OH)2 slurry solution, the total pressure decreased and additional CO2

181 gas was added to the reactor. CO2 gas was continued to be added to the reactor until a desired pressure was reached. This typically took about a day. After an equilibrium condition was established, the lecture bottle was disconnected from the reactor.

The total amount of CO2 gas transferred to the reactor was measured and estimated by the weight and final pressure of the lecture bottle after the experiment.

5.2.2 Calculation Procedure

tot  The total number of moles of CO2 transferred to the reactor ( nCO2 ) was measured by the weight difference in the lecture bottle before and after the experiment.

 The CO2 partial pressure was estimated by taking the difference between the total system tot v tot v pressure ( P ) and water vapor pressure ( PH 2O ) at the temperature ( PCO2  P  PH 2O ).

 The number of moles of CO2 in the vapor phase was estimated using the ideal gas law. v ( nCO2  PCO2V /(RT ) ).

 The moles of dissolved CO2 in the liquid phase was calculated by taking the difference between the total number of moles of CO2 transferred to the reactor and the number of l tot v 182 moles of CO2 present in the vapor phase (i.e. nCO2  nCO2  nCO2 ).

5.3 Results and Discussion

5.3.1 Room Temperature Data

Experimental data were obtained at room temperature using supernatant Mg(OH)2 solution, 0.1 M, and 1 M Mg(OH)2 solutions as summarized in Tables 10-12 and Figure 50.

Table 10. CO2 equilibrium data for supernatant Mg(OH)2 solution

CO2 added PCO2 (kPa) CO2 in the CO2 in the CO2 Mole (mole) vapor phase liquid phase fraction in the (mole) (mole) liquid phase

(xCO2)

0.0434 77.85 0.0261 0.0172 0.0012

0.0981 215.65 0.0725 0.0256 0.0018

0.1522 346.56 0.1165 0.0356 0.0025

0.2295 484.36 0.1629 0.0666 0.0047

0.2925 622.16 0.2093 0.0831 0.0059

The solution obtained at the end of the experiment was very clear indicating all the magnesium hydroxide had been consumed completely. The distribution of CO2 between gas and liquid phases also show that the CO2 is primarily in the gas phase and gradually the amount of

CO2 in the gas phase increases with the pressure. Deionized water was tested at the same pressure and the VLSE data followed the same pattern, which suggests the presence of the slurry is essential for high performance CO2 absorption.

Table 11. CO2 equilibrium data for 0.1 M Mg(OH)2 solution

CO2 added PCO2 (kPa) CO2 in the CO2 in the CO2- Mole (mole) vapor phase liquid phase fraction in the (mole) (mole) liquid phase

(xCO2)

0.151 153.64 0.0516 0.0999 0.0071

0.254 263.88 0.0087 0.1657 0.0117

0.294 408.57 0.1374 0.1573 0.0112

It can be seen that the number of moles of CO2 in the liquid phase reaches a maximum at a certain pressure i.e. [around 40 psia/275 kPa] and after which even if the pressure rises, the amount of CO2 that can be stored in the liquid reaches a plateau. This is best illustrated by figure

50. Also with using slurry, the amount of available magnesium increases and the pH of the sample varied from 6.45 to 6.74. The solution obtained at the end of experiments was slightly yellow with blackish solids which were the impurities present in Mg(OH)2.

Table 12. CO2 equilibrium data for 1 M Mg(OH)2 solution

CO2 added PCO2 (kPa) CO2 in the CO2 in the CO2 Mole (mole) vapor phase liquid phase fraction in the (mole) (mole) liquid phase

(xCO2)

0.453 174.31 0.0586 0.3950 0.0276

0.513 312.11 0.1049 0.4081 0.0285

The 1 M solution can hold a lot more CO2 in the liquid phase than 0.1 M and it has a limiting pressure of somewhere around 30 psia/206 kPa. At the end of the experiment, the residue was found to be milky carbonated slurry.

650 1M Mg(OH) 2 Supernantant Mg(OH) 600 2 0.1 M Mg(OH) 550 2 500 450 400 350 300 250 200

partial pressure (kPa) pressure partial

2 150 100

CO 50 0.000 0.005 0.010 0.015 0.020 0.025 0.030 x (mole fraction in the liquid phase) CO2

Figure 50. VLSE data for Mg(OH)2-CO2-H2O at room temperatures.

From above graph we come to an understanding about the limiting pressure needed for maximum absorption of CO2 by the magnesium hydroxide solution. For the supernatant liquid,

90 higher the system pressure better is the absorption. For 0.1 M slurry solution, the optimum pressure is somewhere around 40 psi. 1 M Mg(OH)2 solution showed a higher CO2 equilibrium capacity than 0.1 M Mg(OH)2 solution.

5.3.2 VLSE Data at 40 oC and 52 oC

o o Experimental data was obtained at 40 C and 52 C using 0.1M and 1M Mg(OH)2 solutions and the data was summarized in the following tables and figures :

o Table 13. CO2 equilibrium data for 0.1 M Mg(OH)2 solution at 40 C

CO2 added PCO2 (kPa) CO2 in the CO2 in the CO2- Mole (mole) vapor phase liquid phase fraction in the (mole) (mole) liquid phase

(xCO2)

0.0972 130.864 0.0396 0.0576 0.00413

0.1127 168.785 0.0510 0.0617 0.00443

0.1454 241.180 0.0729 0.0724 0.00519

o Table 14. CO2 equilibrium data for 0.1 M Mg(OH)2 solution at 52 C

CO2 added PCO2 (kPa) CO2 in the CO2 in the CO2- Mole (mole) vapor phase liquid phase fraction in the (mole) (mole) liquid phase

(xCO2)

0.0970 125.70 0.0386 0.0584 0.00419

0.1281 215.27 0.0661 0.0620 0.00444

0.1413 246.27 0.0756 0.0656 0.00471

300 0.1 M at 52 oC 275 0.1 M at 40 oC

250

225

200

175

Partial Pressure (kPa) Pressure Partial 150

CO 125

0.0030 0.0035 0.0040 0.0045 0.0050 0.0055 x (mole fraction in the liquid phase) CO2

Figure 51. VLSE data for Mg(OH)2-CO2-H2O at higher temperatures.

The solution obtained at the end of the experiment is a clear bubbly liquid. This is a consequence of the low pH which is created due the high operating pressure. The influence of pressure forces the magnesium ions dissociate and go into the solution as magnesium bicarbonate. A commonly observed trend in VLE measurement with amine based solutions is that with an increase in temperature, the equilibrium shifts more towards the vapor phase. This trend is observed in the plot of CO2 partial pressure (kPa) vs. xCO2 (mole fraction in the liquid phase) at 40 oC and 52 oC. The mole fraction values in the liquid phase indicate if the system is operated at higher temperature than 52 oC, it would result in a reduction in the amount of absorbed CO2 in the liquid phase.

o Table 15. CO2 equilibrium data for 1 M Mg(OH)2 solution at 52 C

CO2 added PCO2 (kPa) CO2 in the CO2 in the CO2- Mole (mole) vapor phase liquid phase fraction in the (mole) (mole) liquid phase

(xCO2)

0.2606 113.29 0.0348 0.2258 0.01601

0.2984 187.71 0.0576 0.2407 0.01704

0.3415 253.17 0.0777 0.2638 0.01865

1M Mg(OH) at 52 oC 2 0.1 M Mg(OH) at 52 oC 270 2

240

210

180

150

Partial Pressure (kPa) Pressure Partial

2 120

0.003 0.006 0.009 0.012 0.015 0.018 0.021 x (mole fraction in the liquid phase) CO2

o Figure 52. VLSE data for Mg(OH)2-CO2-H2O at 52 C and different concentrations.

The increase in the concentration of the solution from 0.1 M to 1M at 52 oC resulted in a fourfold increase in the amount of CO2 captured at comparable pressures. This validates the usage of high magnesium hydroxide concentrations in the absorber for maximum sustained CO2 removal.

350

300

250

200

150

100 1M Mg(OH) at 52 oC 2 30 % mass AMP at 40oC

Partial Pressure (kPa) Pressure Partial 2 30 % mass MEA at 40oC 50 0.8 M Piperazine at 55oC o CO 1M MEA at 40 C 50 % MDEA at 55 oC 0 0.2 0.3 0.4 0.5 0.6 0.7 0.8 0.9 1.0 1.1 1.2 1.3 1.4 CO loading (moles of CO /moles of solvent) 2 2

183 168 178 171 Figure 53. Comparison of CO2 loading capacities of various solvents.

Vapor-liquid equilibrium (VLE) data in the case of CO2 capture solvents are generally represented in the form of CO2 partial pressure vs. CO2 loading (moles of CO2/moles of solvent).

CO2 loading values represent the capacity of the solvent and are directly proportional to the absorption capacity of the solvent. Figure compares VLE data available in literature for well- known solvents like MEA [Monoethanolamine], MDEA [Monodiethanolamine], AMP [2-amino-

2-methyl-1- propanol] and piperazine at elevated temperatures, with the experimental values

o obtained for 1M magnesium hydroxide at 52 C. Magnesium hydroxide has a superior CO2 loading capacity compared to the other solvents as demonstrated by this graph. CO2 loading capacities for most solvents decrease with temperature and increase with pressure.

Monoethanolamine at different concentrations show lesser CO2 loading capacities than 1M

o Mg(OH)2 at 52 C. Magnesium hydroxide exhibited a 40 % increase in CO2 loading compared to 50 % MDEA at 55 oC.

o Piperazine [0.8 M] displayed comparable performance with Mg(OH)2 at 52-55 C, but the high cost of piperazine limits its application as a standalone solvent.

5.4 Conclusion

It is highly essential to have equilibrium solubility data over various temperatures, pressures and concentrations for efficient design of the treating units. A static synthetic method was used to determine the vapor-liquid-solid equilibrium data for the CO2–water–

Mg(OH)2system. Experiments were conducted at room temperature for Mg(OH)2 solutions at various concentrations and the need to have a substantial amount of slurry for sustained absorption was established. Experiments were performed at higher temperatures (40 oC & 52 oC) and the shift in equilibrium towards the vapor phase was observed. The CO2 loading capacities

o of Mg(OH)2 at the proposed absorption temperature of 52 C were compared to the values found in literature for various solvents like MEA, MDEA and piperazine at standard operating temperatures respectively. Magnesium hydroxide at 1M at 52 oC was found to give excellent

CO2 loading capacities close to 1 (mol of CO2/ mol of solvent) compared to 0.6 (mol of CO2/ mol of solvent) for 30 % MEA at 40 oC, which is a standard amine concentration used as a bench mark for absorption.

CHAPTER 6. CONCLUSION

Three important phenomena influencing CO2 capture using magnesium hydroxide solutions were investigated, namely: 1) Dissolution rate of magnesium hydroxide, 2) Kinetics of magnesium carbonate formation 3) VLSE of the Mg(OH)2-H2O-CO2 system. Dissolution rate flux was determined using a pH stat apparatus at different operating conditions. An increase in the operational pH decreases the dissolution rate of magnesium hydroxide. The variation of dissolution rate with an increase in temperature was found to follow an Arrhenius type expression, which was used to develop an empirical expression to predict the intrinsic dissolution rate. The dissolution process was modeled using the shrinking core model and surface chemical reaction was found to be the rate controlling step. The apparent reaction rate constant was

-1 -1 o estimated to be 1.1×10 min for an operating pH of 8.6 and temperature of 52 C. Mg(OH)2 was found to exhibit a fractional rate order which are in the range of (0.20

o value at a pH = 8.6 and temperature = 52 C. Regular monitoring of the Mg(OH)2 slurry for

96 estimating particle sizes is recommended to ensure that dissolution takes place at the intrinsic rate, which is essential for sustained CO2 absorption. MINEQL results indicated that operating the absorber at pH of 8.6 would be beneficial to prevent the formation of magnesium carbonate, which is a detrimental byproduct for continuous operation of the slurry. Kinetic analysis of

MgCO3 formation was performed using the mass balance method with solid phase carbon analysis by TGA-MS and liquid phase carbon analysis by HCl titration. The results indicate operating the system at a pH of around 8.6 favors the formation of magnesium bicarbonate, which is helpful for regeneration of the slurry in the stripper phase. Vapor-liquid-solid equilibrium studies were performed by using a static analytical device for the CO2-H2O-

Mg(OH)2 system. The results validate the importance of the slurry in the absorption process and higher the slurry used, better is the performance. Experiments performed at higher temperatures

o o (40 C and 52 C) indicate the shift in equilibrium towards the vapor phase. Mg(OH)2 gives

o better CO2 loading at 52 C compared to amines and other solvents at similar temperatures.

Further research in pilot scale CO2 absorption in packed/bubble columns is being carried out, which will demonstrate the applicability of this system on a large scale.

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