Coal Gasification and Its Applications

COAL GASIFICATION AND ITS APPLICATIONS DAVID A BELL BRIAN F TOWLER MAOHONG FAN Amsterdam • Boston • Heidelberg • Lond...

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COAL GASIFICATION AND ITS APPLICATIONS

DAVID A BELL BRIAN F TOWLER MAOHONG FAN

Amsterdam • Boston • Heidelberg • London • New York • Oxford Paris • San Diego • San Francisco • Singapore • Sydney • Tokyo William Andrew is an imprint of Elsevier

William Andrew is an imprint of Elsevier The Boulevard, Langford Lane, Kidlington, Oxford OX5 1GB, UK 30 Corporate Drive, Suite 400, Burlington, MA 01803, USA First edition 2011 Copyright Ó 2011 Elsevier Inc. All rights reserved. No part of this publication may be reproduced, stored in a retrieval system or transmitted in any form or by any means electronic, mechanical, photocopying, recording or otherwise without the prior written permission of the publisher Permissions may be sought directly from Elsevier’s Science & Technology Rights Department in Oxford, UK: phone (+44) (0) 1865 843830; fax (+44) (0) 1865 853333; email: permissions@elsevier. com. Alternatively you can submit your request online by visiting the Elsevier web site at http:// elsevier.com/locate/permissions, and selecting Obtaining permission to use Elsevier material Notice No responsibility is assumed by the publisher for any injury and/or damage to persons or property as a matter of products liability, negligence or otherwise, or from any use or operation of any methods, products, instructions or ideas contained in the material herein. Because of rapid advances in the medical sciences, in particular, independent verification of diagnoses and drug dosages should be made British Library Cataloguing in Publication Data A catalogue record for this book is available from the British Library Library of Congress Cataloging-in-Publication Data A catalog record for this book is available from the Library of Congress ISBN: 978-0-8155-2049-8 For information on all William Andrew publications visit our web site at books.elsevier.com Printed and bound in Great Britain 11 12 13 14 15 10 9 8 7 6 5 4 3 2 1

INTRODUCTION This is a book about coal gasification and its related technologies. The relationship between these technologies is shown in Figure 0.1. The gasification process begins with a viable feedstock. In this book, we focus on one of those feedstocks that must go through the gasification process, coal. The nature of coal, including its properties and availability, are described in Chapter 1. Petcoke, petroleum coke, a solid, high-carbon byproduct of petroleum refining, can also be gasified. Gasifiers designed for coal, especially high temperature, entrained flow gasifiers, are used for this application. Biomass gasification has a great deal in common with coal gasification, but biomass gasifiers are optimized for biomass feedstock. The product of gasification is syngas, which is primarily a mixture of carbon monoxide and hydrogen. Most syngas, however, is not currently made by gasification, but rather by the steam reforming of natural gas. In this process, steam and natural gas are fed to catalyst-packed tubes, which are held inside a furnace to provide the endothermic heat of reaction. Figure 0.1 also shows other gases, which can be blended with syngas for further processing. One such gas under consideration is hydrogen, which can be produced by electrolyzing water using off-peak power from a nuclear power plant. In a few cases, carbon dioxide from an external source may supplement the carbon monoxide in syngas. Just as coal is not the only feedstock for gasification, gasification is not the only use of coal. Most coal is burned to produce electric power. Chapter 2 describes a few of the non-gasification uses of coal. Gasification is described in Chapters 3, 4, and 5. Chapter 3 describes gasification as a chemical reaction system. Although this chapter may look complex, our knowledge of the chemistry of gasification is far from complete. Chapter 4 covers several gasifier designs. These designs were selected because they are now in commercial use or development, or because they illustrate interesting concepts. One gasification approach is sufficiently different that it deserves its own chapter, underground coal gasification, covered in Chapter 5. Instead of mining coal and transporting it to a gasifier, the coal is left in place underground, and the reactant gases are brought to the coal. Deeply buried coal seams, which are uneconomic to mine, may be exploited by underground coal gasification. Syngas leaving the gasifier contains numerous impurities. The inorganic fraction of the feedstock leaves as solid ash or molten slag. Ash or slag removal is usually an integral part of the gasifier design. If the gasification occurs at relatively low temperatures, then tar will be produced. Tar removal is also an integral part of gasifier design. Highertemperature gasifiers do not produce significant tar. The syngas also contains sulfur in the

ix

x

Introduction

hydrogen, electric power

ammonia, nitrogen fertilizers

methanol, dimethyl ether, hydrocarbons

substitute natural gas, Fischer-Tropsch hydrocarbons

Products

sequestration

steam

Syngas processing

water gas shift

impurity removal

steam reforming

gasification

Gasification

Feedstocks

CO2 removal

coal

petcoke

biomass

natural gas

other gas

Figure 0.1 Gasiﬁcation and related technologies.

form of H2S, with lesser quantities of COS. Sulfur must be removed from syngas either to prevent emission of SO2 when syngas is burned, or to prevent catalyst poisoning in downstream reactors. Sulfur removal is described in Chapter 6. Carbon dioxide removal can occur either as a part of impurity removal, or after water gas shift, as shown in Figure 0.1. The traditional carbon dioxide removal techniques are closely related to sulfur removal, and are described in Chapter 6. The ability to remove carbon from syngas and sequester it in a geological formation is one of the major attractions of coal gasification. This allows coal to be used while minimizing greenhouse gas emissions. A major objection to this approach is that carbon capture and sequestration are expensive. This prompted a great deal of research into new carbon dioxide separation technologies, and which is described in Chapter 10. Syngas contains a number of minor impurities, and one of the more significant is mercury, a neurotoxin. Removal of mercury is discussed in Chapter 9. For some applications, a nearly pure hydrogen stream is desired. In others, such as methanol synthesis, a specific ratio of carbon monoxide to hydrogen is required. In either

Introduction

case, the gasifier usually produces a higher ratio of carbon monoxide to hydrogen than desired. This ratio needs to be shifted towards a greater hydrogen content. The usual way to do this is through the water gas shift reaction in which carbon monoxide reacts with steam to form hydrogen and carbon dioxide, as described in Chapter 7. Hydrogen can then be burned in a turbine to generate electric power, an application known as integrated gasification combined cycle. This is a means of producing electric power from coal with minimal greenhouse gas emissions. Hydrogen is also a potential transportation fuel. The usual approach is to produce electric power from hydrogen in a fuel cell, and then use that power in an electric motor. One of the main technical obstacles is a practical means of storing hydrogen in a vehicle. Chapter 9 explores hydrogen storage for this application. Nearly all synthetic nitrogen chemicals start as ammonia, synthesized from hydrogen and nitrogen gas. Nitrogen fertilizers are, by far, the largest volume synthetic nitrogen chemicals. Chapter 11 describes ammonia synthesis and some of the more common nitrogen fertilizer compounds. Methanol is a major commodity chemical made from syngas, as described in Chapter 12. Methanol is an intermediate used to make a wide range of products. One of these, dimethyl ether (DME), is especially interesting. DME can be used as a fuel or converted to hydrocarbons, including gasoline and olefins for polymer production. Chapter 13 describes the direct conversion of syngas to hydrocarbons, including substitute natural gas (methane) and Fischer-Tropsch liquid, a synthetic crude oil. The Fischer-Tropsch liquid is then refined to meet petroleum product specifications. Coal is an inexpensive feedstock, but gasification-based plants tend to have very high capital construction costs. In concept, one could build a single plant that would incorporate all of the elements shown in Figure 0.1, but such a complex plant would be extraordinarily expensive to build. Instead, gasification-based plants have a more limited set of features dictated by economics and the regulatory environment. There are two major trends that prompt current interest in coal gasification. The first is the widely held belief that conventional petroleum supplies are declining, while demand for transportation fuels continues to rise. This has led to heightened interest in alternative energy supplies, including coal. The second major trend is concern about global warming. Gasification offers a relatively cost-effective means of using coal while minimizing greenhouse gas emissions.

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CHAPTER

1

The Nature of Coal Contents The Geologic Origin of Coal Coal Analysis and Classiﬁcation Coal Rank Ash Thermal Properties Coal as a Porous Material Spontaneous Combustion Reserves, Resources, and Production References

1 2 4 5 9 10 11 15

THE GEOLOGIC ORIGIN OF COAL Coal is fossilized peat. A peat bog is a marsh with lush vegetation. Plant matter dies and falls into the water, where partial decomposition occurs. Aerobic bacteria deplete the water of oxygen, and bacterial metabolic products inhibit further decomposition by anaerobic bacteria. Plant matter accumulates on the marsh bottom faster than it decomposes, and, over a period of many years, a layer of peat forms. The peat that became today’s coal was laid down millions of years ago. Buried peat is converted to coal when high pressure and elevated temperature is applied to the buried layer. This process is known as coalification. The physical and chemical structure of the coal changes over time. As shown in Figure 1.1, the youngest (least converted) coal is known as lignite, which can be further converted to sub-bituminous coal, bituminous coal, and finally anthracite. These coal types strongly influence the properties and use of coal, and will be discussed further.

Peat Lignite

Increasing age, conversion

Sub-bituminous Bituminous Anthracite

Figure 1.1 Coaliﬁcation. Coal Gasification and Its Applications. ISBN B978-0-8155-2049-8.10001-4, doi:10.1016/B978-0-8155-2049-8.10001-4

Ó 2011 Elsevier Inc. All rights reserved.

1

2

The Nature of Coal

Petrography is the visual inspection of a rock sample to determine the mineral types in the sample. When applied to coal, the different coal types are known as macerals. Table 1.1 lists coal macerals, and shows how they are derived from plant material.

COAL ANALYSIS AND CLASSIFICATION Coal is used primarily as a fuel, so its most important property is its heat of combustion. Gross calorific value, also known as higher heating value (HHV), is determined by measuring the heat released when coal is burned in a constant-volume calorimeter, with an intitial oxygen pressure of 2 to 4 MPA, and when the combustion products are cooled to a final temperature between 20 and 35 C (ASTM D 5865-04). The tests mentioned in this book are primarily based on the American Society for Testing and Materials (ASTM) specifications.1 Coal is a variable, widely distributed and widely used material so a wide range of standard tests have been developed by a variety of individuals and organizations. Coal is a porous medium, and these pores, especially in low rank coals, can contain substantial quantities of water even though the coal appears to be dry. The water is either adsorbed onto hydrophilic surface sites or held in pores by capillary forces. When this moist coal is burned or gasified, a substantial fraction of the combustion heat is required Table 1.1 Coal macerals, based on ASTM D121-05 and ASTM D 2799-05a.1 Maceral group Maceral Origin Comments

Vitrinite

Vitrinite

Liptinite

Alginite Cutinite

Intertite

Resinite Sporinite Fusinite Inertodentrinite Macranite Micranite Funginite Secretinite Semifusinite

Woody tissue of plants (cellulose, lignin) Botryoccus algae Waxy coating (cuticle) of leaves, roots and stems Plant resins Spores and pollen grains Some structures of plant cell wall still visible Fragments incorporated within other macerals. No plant cell wall structure, larger than 10 mm No plant cell wall structure, less than 10 mm, and typically 1 to 5 mm Fungi No obvious plant structure, sometimes containing fractures, slits or notch. Like fusinite, but with less distinct evidence of cellular structure.

Most common maceral Waxy, resinous materials

Derived from strongly altered and degraded peat

The Nature of Coal

to vaporize water. Since the final temperature in the gross calorific value test is 20 to 35 C, most of the water is condensed, thereby recovering the heat of vaporization. Water in the HHV test is primarily a non-combustible diluent. For example, a Wyoming Powder River Basin coal typically has an HHV of 19.8 MJ/kg (8500 Btu/lb) and a 28% moisture level. One can then calculate an HHV value for the coal if it is dried: 19:8 MJ=kg MJ Btu ¼ 27:5 11; 800 Eqn. 1.1 HHV ; dry ¼ 1 0:28 kg lb If coal is burned or gasified near atmospheric pressure, then the heat of condensation for the water may not be recovered. For example, in a coal-fired power plant, the water contained in the coal may go up the stack as steam. In other situations, the heat of condensation is recovered, but the value of this heat is relatively low because of its temperature. In these cases, a better estimate of coal heat of combustion is the net calorific value, also known as Lower Heating Value (LHV), which assumes that vaporized water remains as steam and that the heat of condensation is not recovered. Water in the coal reduces its heating value by its heat of vaporization, 2.395 MJ/kg water (1055 Btu/lb water). Again, for a typical PRB coal: 19:8 MJ MJ kg water 2:395 0:28 kg coal kg water kg coal MJ Btu 8; 200 ¼ 19:1 kg coal lb coal

LHV ; moist ¼

Eqn. 1.2

Proximate Analysis (ASTM D 3172-89) involves a series of tests that heat and burn coal. Moisture is measured (ASTM D 3173-03) by determining the weight loss after coal is dried at 104 to 110 C. Volatiles are then measured (ASTM D 3175-02) by determining additional weight loss when coal is pyrolyzed at 950 C. Ash is determined (ASTM D 3174-04) by the weight of inorganic materials remaining after coal is burned. Fixed carbon is the fraction of coal that is not moisture, volatiles, or ash. Fixed carbon, which is mostly carbon but can contain other elements represents the combustible portion of the coal char that remains after the volatiles have been removed. Proximate analysis results are sometimes reported on a dry mineral matter-free basis. Mineral matter is calculated using the following equation: Mm ¼ 1:08A þ 0:55S Where: Mm ¼ percent mineral matter

A ¼ percent ash S ¼ percent sulfur (ASTM D 3177 or D 4239)

Eqn. 1.3

3

4

The Nature of Coal

The 1.08 factor presumes that minerals in the coal are hydrated. This water of hydration is lost when the coal is burned. The 0.55 factor assumes that sulfur is present as pyrites, which in many areas are converted to the corresponding oxides during combustion. Ultimate analysis (ASTM D 3176) describes coal in terms of its elemental composition. For a dried coal, weight percentages of carbon, hydrogen, nitrogen, sulfur, and ash are measured. The remainder of the coal sample is assumed to be oxygen.

COAL RANK In the coalification process, the coal rank increases from lignite to anthracite, as shown in Figure 1.1. Coal rank is useful in the market, because it is a quick and convenient way to describe coal without a detailed analysis sheet. A more detailed description of coal rank is shown in Tables 1.2 and 1.3. Bituminous and sub-bitumous coals are the primary commercial coals. A relatively small amount of anthracite is available. In the USA, anthracites are produced only in northeastern Pennsylvania. Lignites are abundant. But the economics of hauling a low-grade fuel long distances are unfavorable; so most lignite is consumed close to where it is mined. Peat is also mined and generally used close to where it is mined. Peat may be either considered old biomass or very young coal. In nations that regulate greenhouse gas emissions, the difference between the two is more than mere semantics. Carbon dioxide emissions from biomass combustion are not considered a contributor to global warming, because these emissions are offset by carbon dioxide uptake by growing biomass. On the other hand, the same emissions from fossil fuels, are restricted. Emissions from peat combustion are a regulatory gray area. Some coal, particularly bituminous coal, has the tendency to cake. With increasing temperature, coal particles simultaneously pyrolize and partially melt, causing the coal particles to stick to one another. Some gasification reactors, especially moving bed and fluidized bed gasifiers, are limited to processing coal that does not cake. Table 1.2 Classiﬁcation of anthracitic and bituminous coals by rank (ASTM D 388-05).1 Volatile matter limits (dry mineral-matter-free Fixed carbon limits basis), % (dry mineral-matter-free basis), % Equal or Less Greater Equal or greater than than than less than Rank

Meta-anthracite Anthracite Semi-anthracite Low volatile bituminous coal Medium volatile bituminous coal High volatile A bituminous coal

98 92 86 78 69 n/a

n/a 98 92 86 78 69

n/a 2 8 14 22 31

2 8 14 22 31 n/a

The Nature of Coal

Table 1.3 Classiﬁcation of bituminous, sub-bituminous and lignite coals by rank. (ASTM D 388-05). Note that high volatile A bituminous coal is the only rank that is listed in both Table 1.2 and Table 1.3. Gross caloriﬁc value limits (moist, mineral-matter-free basis) Btu/lb MJ/kg Equal or Less Equal or Less Rank greater than than greater than than

High volatile A bituminous coal High volatile B bituminous coal High volatile C bituminous coal Sub-bituminous A coal Sub-bituminous B coal Sub-bituminous C coal Lignite A Lignite B

14 000 13 000 11 500 10 500 9 500 8 300 6 300 n/a

n/a 14 000 13 000 11 500 10 500 9 500 8 300 6 300

32.557 30.232 26.743 24.418 22.09 19.30 14.65 n/a

n/a 32.557 30.232 26.743 24.418 22.09 19.30 14.65

ASH THERMAL PROPERTIES The melting temperatures of coal ash impose temperature limits for coal gasification. Fluidized bed gasifiers and dry-bottom moving bed gasifiers, such as the Lurgi gasifier, require free-flowing ash. The maximum operating temperature for these gasifiers is the initial deformation temperature. When the temperature rises above the initial deformation temperature the ash becomes sticky. Fluidized bed gasifiers often run near the initial deformation temperature to maximize carbon conversion. Entrained flow gasifiers and slagging moving bed gasifiers such as the BGL gasifier require a fluid slag, so they must operate at a sufficiently high temperature to completely melt the ash. Operation at significantly higher temperatures increases oxygen consumption. Ash is a complex mixture of minerals, which will cause the coal ash to melt over a temperature range rather than at a fixed temperature. Temperatures in this range are specified by ASTM D-1857-04. A coal ash cone, 19 mm high and with an equilateral triangle base 6.4 mm on each side, is placed in an oven. Temperatures are reported for reducing or oxidizing gas environments. The initial deformation temperature (IDT) occurs when rounding of the cone tip first occurs. The softening temperature (ST) occurs when the cone has fused to produce a lump which has a height equal to its base. The hemispherical temperature (HT) occurs when the lump height is half the length of its base. The fluid temperature occurs when the fused mass has spread out in a nearly flat layer with a maximum height of 1.6 mm. A number of researchers have attempted to correlate ash thermal properties with ash composition. The most extensive effort was by Seggiani and Pannocchia,2 who correlated the behavior of 433 ash samples, based on nine elemental concentrations.

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6

The Nature of Coal

Note that mineral elemental compositions are reported as if the mineral sample were a blend of simple metal oxides. For example, the fraction of aluminum in a sample is typically reported as the equivalent weight percent of Al2O3. Seggiani and Pannocchia’s correlations are based on mole percents, rather than weight percents, on a normalized, SO3-free basis. The correlation for initial deformation temperature is given as: IDT ; C ¼ 2; 040 exp 0:1

SiO2 SiO2 þ Fe2 O3 þ CaO þ MgO

2

þ 83:4P2 O5

þ 2:12 Al2 O3 þ 39:3TiO2 þ 0:335 ðFe2 O3 Þ2 þ 0:118 ðAl2 O3 Þ2 þ 0:135 ðCaOÞ2 0:116 ðSiO2 Þ ðFe2 O3 Þ þ 0:0768 ðSiO2 Þ ðAl2 O3 Þ SiO2 2 þ 0:533 ðFe2 O3 Þ ðCaOÞ þ 2:42 Al2 O3 2 SiO2 þ 205 SiO2 þ Fe2 O3 þ CaO þ MgO 4 þ 780 exp 10 ðSiO2 Þ ðAl2 O3 Þ 2170 Eqn. 1.4 The correlation for softening temperature is given as: 2 SiO2 ST ; C ¼ 5; 360 exp 0:1 þ 91:3 P2 O5 SiO2 þ Fe2 O3 þ CaO þ MgO þ 0:282 ðFe2 O3 Þ2 þ 0:178 ðCaOÞ2 þ 0:939ðMgOÞ2 þ 0:630 ðFe2 O3 Þ SiO2 2 ðCaOÞ 1:03 ðFe2 O3 Þ ðMgOÞ þ 2:34 Al2 O3 2 SiO2 140 SiO2 þ Fe2 O3 þ CaO þ MgO CaO þ MgO 85:9 Fe2 O3 þ CaO þ MgO þ K2 O þ Na2 O3 þ 3; 120 exp 104 ðSiO2 Þ ðAl2 O3 Þ 7820 Eqn. 1.5

The Nature of Coal

The correlation for hemispherical temperature is given as: HT ; C ¼ 2; 150 exp 0:1

SiO2 SiO2 þ Fe2 O3 þ CaO þ MgO

2

þ 53:1 TiO2 Fe2 O3 2 2 25:3K2 O þ 16:0 ðTiO2 Þ þ 0:0877 ðCaOÞ þ 19:3 CaO 2 SiO2 þ 0:285 Al2 O3 2 Fe2 O3 þ CaO þ MgO þ K2 O þ Na2 O þ 910exp 0:1 1 SiO2 þ Al2 O3 þ TiO2 þ P2 O5 Fe2 O3 þ CaO þ MgO þ K2 O þ Na2 O 2 þ 41:9 SiO2 þ Al2 O3 þ TiO2 þ P2 O5 2 Fe2 O3 þ CaO þ MgO þ K2 O þ Na2 O þ 86:4 1 SiO2 þ Al2 O3 þ TiO2 þ P2 O5 2 SiO2 þ 216 2; 120 SiO2 þ Fe2 O3 þ CaO þ MgO Eqn 1.6

The correlation for fluid temperature is given as: 2 SiO2 þ 6:13Al2 O3 FT ; C ¼ 2; 240 exp 0:1 SiO2 þ Fe2 O3 þ CaO þ MgO

þ 58:0TiO2 13:8MgO þ 0:259 ðFe2 O3 Þ2 þ 0:278 ðAl2 O3 Þ2 þ 0:736 ðMgOÞ2 þ 0:259 ðFe2 O3 Þ ðCaOÞ 0:730 ðFe2 O3 Þ ðMgOÞ SiO2 2 Fe2 O3 þ CaO þ MgO þ K2 O þ Na2 O 2 þ 92:0 þ 2:03 Al2 O3 SiO2 þ Al2 O3 þ TiO2 þ P2 O5 2 SiO2 þ 231 1; 340 SiO2 þ Fe2 O3 þ CaO þ MgO Eqn. 1.7 The temperature of critical viscosity, Tcv, is not part of the ASTM D1857 test but it is important for slagging gasifiers because it marks the transition of slag from a

7

8

The Nature of Coal

difficult-to-handle Bingham plastic, below Tcv, to a more easily handled Newtonian fluid, above Tcv. The correlation for temperature of critical viscosity is given as: Tcv ; C ¼ 935P2 O5 þ 4:11Al2 O3 þ 2; 580ðP2 O5 Þ2 þ 0:254ðAl2 O3 Þ2 0:139ðNa2 OÞ2 þ 0:108 ðSiO2 Þ ðFe2 O3 Þ þ 0:0377 ðSiO2 Þ ðAl2 O3 Þ SiO2 SiO2 2 þ 14:0 þ 0:00691 Fe2 O3 þ CaO þ MgO þ 3:05 Al2 O3 Al2 O3 Fe2 O3 þ CaO þ MgO þ K2 O þ Na2 O þ K2 O þ Na2 O 2 þ 7:40 SiO2 þ Al2 O3 þ TiO2 þ P2 O5 Fe2 O3 þ CaO þ MgO þ K2 O þ Na2 O 2 113 SiO2 þ Al2 O3 þ TiO2 þ P2 O5 Fe2 O3 þ CaO þ MgO þ K2 O þ Na2 O 5:48 ðNa2 OÞ SiO2 þ Al2 O3 þ TiO2 þ P2 O5 Fe2 O3 164 Fe2 O3 þ CaO þ MgO þ K2 O þ Na2 O Fe2 O3 þ CaO þ MgO þ K2 O þ Na2 O 1 7:40 SiO2 þ Al2 O3 þ TiO2 þ P2 O5 þ 409 exp 104 ðSiO2 Þ ðAl2 O3 Þ þ 675 Eqn. 1.8 Seggiani and Pannocchia report standard deviations for their correlations to be 70 to 88oC. Table 1.4 compares experimental results for four American coals from Baxter3 to the temperatures predicted by these correlations. The predicted results are very close to the experimental results for the lignite and the sub-bituminous coals. The exception is the predicted temperatures are substantially higher than the experimental values for the bituminous coals. Inorganic additives have been added to coal gasifiers to modify ash thermal properties. For example; alkaline materials such as sodium, potassium and calcium compounds tend to lower ash melting temperatures. These can be added to an entrained flow gasifier to lower slag viscosity. Care must be taken with refractrory-lined gasifiers, because these compounds may attack the refractory. The opposite approach was taken by van Dyk and Waanders.4 They sought to increase the ash fusion temperature (ISO 540 and 1195E) to allow higher temperature operation in a Lurgi gasifier. Tests with Al2O3, TiO2, and SiO2 showed that Al2O3 was most effective. Addition of 6 weight % Al2O3 boosted the ash fusion temperature of a mixture of South African coals from 1,340 C to greater than 1,600 C.

The Nature of Coal

Table 1.4 Coal ash thermal properties: comparison of experimental values to values predicted by the Seggiani and Pannocchia correlations. Wyodak subPittsburgh No. 8 Illinois No. 6 Coal rank Beulah lignite bituminous bituminous bituminous Mole Weight Mole Weight Mole Weight Mole Weight Mole Mineral weight % % % % % % % %

Al2O3 101.94 CaO 56.08 Fe2O3 159.70 K2O 94.20 MgO 40.32 Na2O 61.99 141.96 P2O5 SiO2 60.06 79.90 TiO2 Total IDT, C, exp. IDT, C, calc. Error, C HT, C, exp. HT, C, calc. Error, C FT, C, exp. FT, C, calc. Error, C

13.968 16.358 12.249 0.220 4.457 6.501 0.001 21.227 0.415 73.395 1,108 1,061 47 1,176 1,177 þ1 1,199 1,220 þ21

12.666 26.963 7.090 0.216 10.219 9.694 0.001 32.671 0.480 100.00

14.218 24.845 5.450 0.198 4.363 0.872 1.009 26.537 1.309 78.801 1,183 1,187 þ4 1,212 1,281 þ69 1,253 1,229 25

11.562 36.725 2.829 0.175 8.970 1.166 0.589 36.627 1.358 100.00

20.657 2.085 29.238 1.742 0.786 0.403 0.1151 41.696 0.896 97.653 1,047 1,189 þ142 1,082 1,400 þ318 1,222 1,467 þ244

17.262 3.167 15.596 1.576 1.660 0.554 0.091 59.140 0.956 100.00

16.904 5.180 20.671 2.031 0.798 1.302 0.167 46.854 0.883 94.790 1,060 1,180 þ120 1,090 1,342 þ251 1,253 1,421 þ168

13.348 7.435 10.419 1.735 1.593 1.691 0.095 62.794 0.890 100.00

Ash is typically land-filled. If the landfill is unlined, then water percolating through the ash pile may affect surface and groundwater quality. Many ashes are alkaline and there is the possibility that toxic heavy metals in the ash may be leached by rainwater. Slagging gasifiers produce glassy, non-leachable slag. Some coal ashes are pozzolanic, which means that they tend to set up like cement when mixed with water. These ashes are often used as road base. High calcium ash has an analysis that is similar to commercial cement. Pozzolanic ashes are less likely to pose leachate problems than unconsolidated ashes.

COAL AS A POROUS MATERIAL Coal is a porous material.5 Pores are classified as macropores (greater than 50 nm), which are measured using mercury porosimetry, mesopores (2.0 to 50 nm), measured by nitrogen adsorption at 77 K, and micropores (0.4 to 2.0 nm), measured by carbon dioxide adsorption at 298 K. Micropores are due to the voids from imperfect packing of large organic molecules. Coals typically have surface areas in the range of 100 to 400 m2/g,

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The Nature of Coal

which is due almost entirely to micropores. For comparison, a typical atomic diameter is 0.25 nm, so only small molecules may penetrate micropores. Olague and Smith6 studied gas diffusion in coal.

SPONTANEOUS COMBUSTION Coal oxidizes when it is exposed to oxygen at ambient conditions. Low grade coals are especially prone to low temperature oxidation. The effect of long-term air exposure on coal quality is known as weathering. Oxidation at low temperatures is exothermic, resulting in increased temperatures that accelerate the rate of coal oxidation. This sometimes leads to spontaneous combustion of coal. Itay et al.7 studied low temperature oxidation of South African coal. They found that the quantity of oxygen adsorbed was greater than the quantity of oxygen-containing product gasses (CO2, CO, H2O), so most of the adsorbed oxygen remains in the coal. Other investigors8,9 found an increase in carboxylic acids in weathered coal. Itay et al. found that oxygen uptake with repeated oxygen exposures declines. This same effect is shown in Figure 1.210. Small particles tend to oxidize faster than large particles, but the particle size effect is not large. This suggests that the rate of oxidation is limited by oxygen diffusion in coal micropores or by surface reaction rates. As-mined low grade coals typically have high water content, and the water-filled pores tend to block low temperature oxidation. As shown in Equation 1.1, the heat content of these coals can be greatly increased by drying. Unfortunately, the dry coals cannot be safely stored or shipped under ambient conditions due to their tendency to spontaneously combust. Exposure of dry coals to high humidity or liquid water

Figure 1.2 Low temperature oxidation of about 1.5 g dried, sub-bituminous, Adaville coal from Kemmerer, Wyoming in a microcalorimeter10. Note that the heat released diminishes with repeated oxygen exposures.

The Nature of Coal

accelerates the rate of low-temperature oxidation, possibly because water adsorption on coal is exothermic. Processes have been developed11,12 to convert high moisture, low grade coals to low moisture fuels with reduced spontaneous combustion tendencies.

RESERVES, RESOURCES, AND PRODUCTION Throughout history coal has played a very small role in the world’s energy mix. Locally it was a curiosity because it was an interesting rock that could be made to burn. However, commencing in about the year 1500, it began to be mined for small scale energy use in England and Germany. When the Industrial Revolution dawned in England in the eighteenth century coal became a significant energy source that fueled the English factories that were the hallmark of the Industrial Revolution. However, in terms of total energy use biomass (particularly wood) remained the major source of energy for the world until about 1900, when coal overtook biomass as the chief world energy source. In the United States, where coal was abundant and easier to mine, it had become the chief energy source in about 1880. Throughout the first half of the twentieth century coal was the major world energy source, until it was overtaken by oil in about 1960. Even though coal production has continued to increase since then it remains in second place behind oil and just ahead of natural gas. This is illustrated in Figure 1.3. In the near future it is likely to remain in the second position until oil production peaks and starts to decline. It is conceivable that when this happens coal will again

Figure 1.3 World energy sources since 1800.

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12

The Nature of Coal

become the most consumed energy source in the world. In terms of energy reserves the world has much more coal than any other energy source. Some might argue that coal production will be restricted because of the amount of CO2 that it produces. But as we learn to capture and sequester the CO2 economically, coal production will likely continue to increase. The entire quantities of coal present, regardless of the cost or practicality of recovery, are known as resources. A 5 cm thick layer of low quality coal buried under 2000 m of overburden contributes to total resources, but it is unlikely that it will ever be mined. A more practical measure of the quantities of coal available are known as reserves, which is the subset of resources that can be mined at current prices using current technology. There are an estimated13 275 billion tons of coal reserves in the United States, compared to approximately 4 trillion tons of coal resources. Table 1.5 shows estimates of worldwide reserves, recent annual production and reserve/production (R/ P) ratios, which is the number of years these reserves will last at current production levels.14

Table 1.5 Worldwide coal production and reserves.14 Region and selected 2006 coal production, 2003 coal reserves, countries million short tons million short tons

United States North America Columbia Central and South America Germany Greece Poland Turkey Europe Kazakhstan Russia Former USSR Iran Middle East South Africa Africa Australia China India Indonesia Asia and Oceania World

Reserves/production, years

1,161.44 1,243.47 70.22 86.55

270,718 279,506 7,287 21,928

233 225 104 253

222.74 72.37 171.12 71.51 799.27 106.17 340.61 534.21 0.94 0.94 269.37 274.93 419.58 2,620.50 497.18 186.25 3,842.16 6,781.53

7,428 4,299 15,432 4,614 65,572 34,479 173,074 250,694 462 462 53,758 55,486 86,531 126,215 101,903 5,476 327,264 1,000,912

33 59 90 65 82 325 508 469 491 491 200 202 206 48 205 29 85 148

The Nature of Coal

It does not mean that the world will run out of coal in the next 148 years. For example, Luppens, et al.15 studied coal reserves and resources in the Gillette Coal Field, a 5,180 km2 portion of the 57,000 km2 Powder River Basin in northeastern Wyoming and southeastern Montana. Coal is abundant throughout the Powder River Basin, but mining is restricted to coal with the lowest mining cost. Most mining in the basin occurs within the Gillette Coal Field, where thick coal seams lie close to the surface. There are multiple coal beds and the thickest of these, the Anderson, has an average 15 m thickness. The beds dip slightly from east to west, so coal is produced from 13 strip mines along a 78 km northesouth line along the eastern edge of the district. These 13 mines produced over 42% of the coal produced in the USA in 2007.14 Because of the abundance of this coal, and because of its low mining cost, the open market mine mouth price14 for this sub-bituminous coal in 2007 was $9.67/ton; compared to $47.63/ton for bituminous coal from West Virginia. As shallow coal in the Gillette Coal field is depleted, mining moves to the west, with gradually increasing overburden and gradually increasing mining cost. When the overburden becomes too thick for strip mining, underground mining may be used to further extract coal. Luppens et al. estimated the volume of coal available in the Gillette Coal Field versus price, and these data are shown in Figure 1.4. At the 2008 production rate of 464 million tons/year, the coal reserves in the Gillette Coal Field will last only 21 years if coal is priced at $10/ ton. If the price of coal rises to $60/ton, then the coal will last 176 years at the 2008 production rate. Of course, the Gillette Coal Field is only one mining district, amongst many scattered throughout the world. As mining costs in the Gillette Coal field rise, mining will shift to other portions of the Powder River Basin and to other coal provinces. On a worldwide basis, coal will be available for a very long time, but coal prices are expected to gradually increase due to increasing mining costs.

$60

Coal Cost, $/ton

$50 $40 $30 $20 $10 $0 0

10,000

20,000

30,000

40,000

50,000

60,000

Coal reserves, millions of tons

Figure 1.4 Coal reserves versus cost in the Gillette Coal Field.15

70,000

80,000

13

14

The Nature of Coal

In-situ coal gasification is the process of partially burning a coal seam in place to produce a synthesis gas. This technology may greatly expand coal reserves, because deeply buried coal seams may be exploited at a reasonable cost. In-situ coal gasification has its own set of issues, however, which will be discussed in Chapter 4. We cannot continue to mine coal until we completely exhaust the resource. Other phenomena will curb mining before then. One concept, popular in peak oil discussions, is called Energy Returned on Energy Input (EROEI). This concept states that if the energy consumed in producing energy is greater than energy produced, then energy production will halt regardless of price. Coal production will probably never reach the EROEI limit. Instead, the limit to coal mining will be set by the principle of economic substitution. Wooly mammoths were once a major human food source. When mammoths were hunted to extinction during the last ice age, our ancestors did not starve. Instead, they found something else to eat. This is an early example of economic substitution. Coal initially became popular when the growing human population and increased urbanization made wood scarce. The technology of coal use developed to the point that coal became a major source of chemicals, fuel gas and transportation fuel. These coal uses fell out of favor when natural gas and crude oil became abundant and inexpensive in the mid twentieth century. Current interest in coal technology is largely the result of rising natural gas and crude oil prices and concern about future energy supplies. At the time this was written, almost all coal consumed in the industrialized nations was burned to produce electric power. In the near future, coal-burning power plants will probably be required to capture and sequester carbon dioxide to reduce global warming. When this happens, coal may no longer be a low cost fuel for power generation. The estimated cost16,17 of pulverized coal-fired power production with carbon capture and sequestration is about double the cost of pulverized power production without carbon capture and sequestration. Residential power customers would see a 50% increase in their power bills, assuming that distribution costs would change. At this price, non-fossil fuel sources of electric power, such as nuclear and wind power, are attractive. Currently, solar power is too expensive, but advancing technology promises to lower costs. A challenge for clean coal technology is to produce power and sequester carbon dioxide at a price that is competitive with alternative power sources. Integrated gasification combined cycle (IGCC) coal plants with carbon capture and sequestration (CCS) have attracted a great deal of interest because the cost of electric power from these plants is only 60% higher than conventional pulverized coal generated electricity.16 The cost and complexity of replacing coal-fired power plants is enormous, so coal-fired power plants will be a substantial source of electric power for many years. Coal has a bright future as a raw material for liquid fuels, fuel gas, and chemicals. Projected crude oil prices suggest that liquid fuels will be produced from coal at a lower

The Nature of Coal

cost than the same fuels produced from crude oil. Coal technology, and in particular, coal gasification, will have a large role in energy production for many years.

REFERENCES 1. American Society for Testing and Materials, Annual Book of ASTM Standards 2006, Volume 5 Part 6, Gaseous Fuels; Coal and Coke. 2. Seggiani M, Pannocchia G. Prediction of coal ash thermal properties using partial least-squares regression. Ind Eng Chem Res. 2003;42:4919-4926. 3. Baxter L. Brigham Young University. Coal database, www.et.byu.edu/%7Elarryb/CoalDatabase.htm. 4. van Dyk JC, Waanders FB. Manipulation of gasification coal feed in order to increase the ash fusion temperature of the coal enabling the gasifiers to operate at higher temperatures. Fuel. 2007;86:27282735. 5. Gan H, Nandi SP, Walker L. Nature of the porosity in American coals. Fuel. 1972;51:272-277. 6. Olague NE, Smith DM. Diffusion of gases in American coals. Fuel. 1989;68:1381-1387. 7. Itay M, Hill CR, Glasser D. A study of the low temperature oxidation of coal. Fuel Proc Tech. 1989;21:81-97. 8. Yun Y, Meuzelaar HLC. Development of a reliable coal oxidation (weathering) indexdslurry pH and its applications. Fuel Proc Tech. 1991;27:179-202. 9. Hayashi J, Aizawa S, Kumagai H, et al. Evaluation of a brown coal by means of oxidative degradation in aqueous phase. Energy & Fuels. 1999;13:69-76. 10. Balasubramani R. Calorimetric Investigation of the kinetics of low-temperature oxidation of dry coal, M.S. Thesis, University of Wyoming (2003). 11. Sethi VK, Dunlop DD. A coal upgrading technology for sub-bituminous and lignite coals, <www. westernresearch.org/uploadedFiles/Energy_and_Environmental_Technology/Coal/Upgrading_ (Including_Headwaters)/CoalUpgradingFMI.pdf>. 12. Evergreen Energy, K-Fuel and K-Direct, <www.evgenergy.com/fact_sheets/K-Fuel.pdf>. 13. U.S. Energy Information Administration, “U.S. coal reserves, 1997 Update,” DOE/EIA-0529(97), 1999. 14. U.S. Energy Information Administration, <www.eia.doe.gov/fuelcoal.html>. 15. Luppens JA, Scott DC, Haacke JE, et al. Assessment of coal geology, resources, and reserves in the Gillette Coalfield, Powder River Basin, Wyoming, U.S. Geological Survey, Open-File Report 20081202 (2008). 16. Katzer J, Ansolabehere S, Beer J, et al. The future of coal, options for a carbon-constrained world, an interdisciplinary MIT study (2007). 17. Woods MC, Capicotto PJ, Haslbeck JL, et al. Cost and performance baseline for fossil energy plants, volume 1: Bituminous coal and natural gas to electricity, DOE/NETL-2007/1281, <www.netl.doe. gov/energy-analyses/pubs/Bituminous%20Baseline_Final%20Report.pdf> (2007).

15

CHAPTER

2

Non-gasiﬁcation Uses of Coal Contents Home Heating and Cooking vs. Industrial Use Coal Combustion Pollutants Pulverized Coal Combustion Supercritical Pulverized Coal Combustion Carbon Capture with Pulverized Coal Combustion Plants Oxy-combustion Sargas Coal-to-liquids ENCOAL Direct Hydrogenation of Coal References

17 17 19 20 21 24 27 28 28 30 33

HOME HEATING AND COOKING VS. INDUSTRIAL USE The simplest use of coal is to burn it for heat. Coal was once used as a household heating and cooking fuel in Western nations; but it was largely replaced by natural gas, propane, electricity and fuel oil. Coal is still used for household heating and cooking in China, where it is a major source of air pollution. In Western nations, coal is used primarily as a fuel for large industrial boilers, especially for electric power generation. Large users are able to get more complete combustion, which reduces odor and soot, and are able to install complex and expensive air pollution equipment. Since pollution control equipment strongly influences the configuration of a modern coal-burning plant, emissions from coal combustion will be described next.

COAL COMBUSTION POLLUTANTS The US Environmental Protection Agency developed a list of priority pollutants, which are common air pollutants that are primarily generated by combustion. The following is a partial list of these pollutants. SOx consists primarily of SO2 but may also contain small amounts of SO3. In the atmosphere, SO2 oxidizes to SO3. This combines with water to form sulfuric acid, H2SO4, the primary acid component in acid rain. Combustion of sulfur-containing fuels creates SOx, and coal typically has high sulfur levels compared to other fossil fuels. Coal Gasification and Its Applications. ISBN B978-0-8155-2049-8.10002-6, doi:10.1016/B978-0-8155-2049-8.10002-6

Ó 2011 Elsevier Inc. All rights reserved.

17

18

Non-gasiﬁcation Uses of Coal

NOx consists of several nitrogen-oxygen compounds that contribute to photochemical smog, ozone depletion and global warming. There are two primary sources of NOx. Fuel NOx forms when nitrogen-containing fuel is burned. Not all nitrogen in the fuel forms NOx. Some of the fuel nitrogen may be converted to N2. Thermal NOx is created by direct combination of N2 and O2 in a flame. Thermal NOx is favored by high flame temperatures and high oxygen concentrations. At ambient conditions, NOx is not thermodynamically stable; but it is very difficult to decompose once formed. CO is formed when carbon-containing fuels are burned. In a flame, carbon is burned to form CO; which is then further oxidized, at a slower reaction rate, to CO2. In nearly all combustion processes, some of the intermediate product, CO, escapes into the flue gas. Carbon monoxide emissions are favored by low oxygen/fuel ratios. Particulates are divided into two categories, PM10, which consists of particles less than 10 microns in diameter; and PM2.5, a subset of PM10 which consists of particles less than 2.5 microns in diameter. When inhaled, these particles, especially PM2.5, tend to remain in the lungs. This can lead to chronic health conditions such as black lung in coal miners, silicosis in people who have prolonged exposure to dust and smoker’s lungs. Some dust is generated when coal is mined, crushed, and shipped. When coal is used for home heating and cooking, the flue gas can contain significant quantities of soot, which is a fine carbon-rich dust. In industrial boilers, combustion is more complete and little soot is produced. Particulate emissions are primarily due to fly ash, which are the fine ash particles entrained in the flue gas. Volatile organic compounds, VOCs, are nearly all organic compounds that have a significant vapor pressure at ambient conditions. In home heating and cooking applications, VOCs in the flue gas cause disagreeable odors. Since combustion is more complete in industrial boilers, little odor is produced. VOCs are a major issue in organic chemical plants, including coal-to-chemical and coal-to-liquid fuels plants. Air toxics include a long list of specific toxic compounds. Coal contains small quantities of volatile heavy metals; which vaporize during combustion and may leave with the flue gas. Mercury1 has received the most attention. Mercury is a neurotoxin, and tends to accumulate in aquatic systems. Mercury bio-accumulates, meaning that large fish that eat smaller mercury-containing fish do not excrete the mercury. Consequently, mercury concentrations are highest at the top of the food chain, including large fish and the people who eat them. The US Environmental Protection Agency issued the first Clean Air Mercury Rule in 2005. Greenhouse gasses include CO2, CH4, and NOx compounds. Because of the large volume emitted, CO2 has received the most attention. Most members of the scientific community believe that global warming is largely due to greenhouse gasses released by fossil fuel combustion. Since coal has lower H/C ratios than other fossil fuels, coal combustion releases more CO2 per unit of energy than other fossil fuels.

Non-gasiﬁcation Uses of Coal

PULVERIZED COAL COMBUSTION The most common type of coal-fired power plant is pulverized coal combustion (PCC), shown in Figure 2.1. A mixture of pulverized coal and air is blown into a low NOx burner. This burner has an annular arrangement. Coal and a portion of the air are fed to the center tube. The remainder of the air is fed through the space between the inside and outside tubes. The main portion of the flame has a low oxygen/fuel ratio and a relatively low temperature, both of which inhibit formation of NOx. The additional air oxidizes CO to CO2. A low NOx burner reduces NOx emissions from about 11 kg to about 5.5 kg per ton of sub-bituminous coal.2 The walls of the furnace have a water wall construction, meaning that side-by-side tubes are welded together to form a continuous wall. Hot combustion gas first rises through the boiler section where pressurized water is boiled to make steam. Next comes the superheater section, where the steam temperature is raised above its boiling point. Then the economizer section preheats the boiler feed water. Finally there is a rotating plate exchanger. Iron plates rotate into the path of the warm flue gas. The warm plates then rotate out of the flue gas path and into the air path, where the plates preheat combustion air. Flue gas then enters the selective catalytic reactor (SCR). Ammonia is injected into the SCR where it reacts with NOx (here shown as NO) to form N2 and H2O. This eliminates 75 to 85% of the NOx.

Superheater section

Economizer section

Boiler section

Air flue gas

Low NOx burner Rotating plate exchanger

Coal + air

Preheated air

air

Bottom ash

Figure 2.1 Pulverized coal combustion plant.

Flue gas to SCR

19

20

Non-gasiﬁcation Uses of Coal

6NO þ 4NH3 /5N2 þ 6H2 O

R-2.1

The flue gas then enters the bag-house, which removes fly ash. The flue gas is forced to flow through a bag filter that captures the fly ash. Some power plants use an electrostatic precipitator in place of a bag-house. In this process the flue gas flows between two electrified parallel plates. These plates attract the fly ash to the surface of the plates where it is held through the electro-static force. The de-ashed flue gas flows on through the plates. A flue gas desulfurization unit (FGD) uses a wet or dry limestone (CaCO3) stream to convert SO2 to gypsum (CaSO4$2H2O), which is land-filled. SO2 ðgÞ þ CaCO3 ðsÞ/CaSO3 ðsÞ þ CO2 ðgÞ

R-2.2

2CaSO3 ðsÞ þ O2 ðgÞ þ 4H2 OðlÞ/2CaSO4$2H2 OðsÞ

R-2.3

At the time this was written, there was no standard method for mercury control. Control techniques under consideration include flotation of the coal to remove mineral matter and injection of activated carbon into the flue gas ahead of the bag-house. Mercury in the flue gas can be in either oxidized or un-oxidized form. Halide salts have been used to converted un-oxidized mercury to mercuric halides, which are more readily removed in the bag-house or FGD. The Western Research Institute has also patented a process where mercury is removed by preheating the coal to a particular temperature prior to combustion.

SUPERCRITICAL PULVERIZED COAL COMBUSTION In a simplistic thermodynamic analysis, a pulverized coal combustion plant may be viewed as a classic heat engine, shown in Figure 2.2. Although a real PCC unit is much more complex than what is shown here, this simplistic picture can be used to illustrate a trend. The theoretical maximum efficiency, hmax, is given by the Carnot cycle: hmax ¼ Heat in

Steam turbine

Heat out

TH TC TC ¼ 1 TH TH

TH = temperature of steam from boiler

work

electricity

TC = temperature of cooling water

Figure 2.2 Pulverized coal combustion plant as a classic heat engine.

Eqn. 2.1

Non-gasiﬁcation Uses of Coal

In theory, one may increase the efficiency by increasing TH or decreasing TC. Most power plants use cooling water from an evaporative cooling tower. Smaller numbers of power plants use cooling water on a once-through basis or use air cooling. The temperature of the water or air used to cool the power plant effectively sets TC. This means that to increase efficiency TH must be increased. The maximum steam temperature and pressure is set by the steam tube materials of construction. Metal strengths fall with increasing temperature, therefore these tubes must resist the corrosive environment in the furnace. Steam tube metallurgy is an active research area. The latest steam tubes allow operation above the critical pressure of water, as shown in Table 2.1. Table 2.1 Steam conditions and efﬁciency for subcritical and supercritical pulverized coal combustion.3 Subcritical pulverized Supercritical pulverized Water critical coal combustion coal combustion point

Boiler pressure, MPa Boiler temperature, C HHV efficiency,%

16.5 566 36.8

24.1 593 39.1

21.94 374

The higher efficiency of the supercritical plant means that less coal is needed to produce the same amount of power. This also reduces the corresponding emissions. This is a small but significant effect.

CARBON CAPTURE WITH PULVERIZED COAL COMBUSTION PLANTS The clean coal concept generally refers to a power plant that burns coal, or a coal-derived fuel such as the syngas produced by a coal gasifier. It then separates the CO2 and sequesters it to prevent emission of CO2 to the atmosphere. Sequestration can take a variety of forms, but the most common approach is to compress CO2 and store it underground. In a pulverized coal combustion plant, the following three steps are required: 1. Separate CO2 from flue gas. 2. Compress CO2, typically to about 15 MPa. 3. Inject CO2 into a porous geologic formation. To illustrate the difficulty of step 1, consider a perfect membrane illustrated in Figure 2.3. This hypothetical membrane has perfect selectivity for CO2 and offers no resistance for CO2 transport. Since the membrane offers no resistance for CO2 transport, the CO2 partial pressure is the same on both sides of the membrane: PCO21 ¼ PCO22

Eqn. 2.2

21

22

Non-gasiﬁcation Uses of Coal

membrane

PCO2-1

PCO2-2

CO2 flue gas

Figure 2.3 Hypothetical perfect membrane for CO2 separation from ﬂue gas.

If the flue gas is at standard atmospheric pressure, 101 kPa and the flue gas contains 13% CO2; then CO2 separation will not begin until the pressure on the CO2 side drops below 13 kPa. Removal of 90% of the CO2 requires a 1.3 kPa CO2 pressure and 99% removal requires a 0.13 kPa CO2 pressure. Even a perfect membrane would require large and expensive vacuum pumps to separate CO2 from the flue gas. Real membranes with less than perfect selectivities and significant transport resistance would be more costly. The usual approach to CO2 removal is to use a liquid absorbent or a solid adsorbent that has an affinity for CO2. This allows the CO2 to be removed at atmospheric pressure. A weak bond is formed between the CO2 and the liquid or solid. This bond is then broken, usually by heating. This will regenerate the absorbent or adsorbent and free the CO2. A strong bond between CO2 and the liquid or solid leads to fast and nearly complete removal of CO2 from flue gas. This strong bond requires a large amount of energy to break; so selection of the absorbent or adsorbent is a compromise between effective CO2 removal and ease of regeneration. Aqueous amine solutions have long been used for removal of CO2 and other acid gasses, such as SO2 and H2S, from gas streams. One of the more common commercial amines is mono-ethanol-amine (MEA), shown in Figure 2.4. HH N-C-C-O-H H HH

H

Figure 2.4 Mono-ethanol-amine (MEA), an absorbent used for CO2 removal.

Figure 2.5 shows a simplified process for the removal of CO2 from power plant flue gas. The warm flue gas is cooled by water evaporation in a direct contact cooler. An aqueous solution of MEA is then used to absorb CO2 from the flue gas. A water wash section above the MEA absorption section removes traces of MEA from the flue gas. The CO2-loaded MEA solution is then sent to a stripper, where CO2 is boiled off the MEA solution. The regenerated MEA solution is cooled and sent to the absorber. The CO2 is compressed and sequestered. Woods et al.3 compared designs for a subcritical pulverized coal combustion power plant with and without carbon capture and sequestration (CCS). For both cases, the feed

Non-gasiﬁcation Uses of Coal

CO2 to compression

CO2-free flue gas to atmosphere

condenser water MEA solution

water

cooler

Stripper

Absorber

water

CO2 – loaded solution

Direct Contact Cooler

reboiler

flue gas water recycle

Figure 2.5 Mono-ethanol-amine (MEA) based process for CO2 removal from ﬂue gas.

was an Illinois No. 6 bituminous coal. The CCS case used an MEA process similar to that shown in Figure 2.5. Table 2.2 shows a comparison of estimated costs and efficiencies for these two plant designs. The carbon capture and sequestration system requires considerable energy, especially for the stripper reboiler heat and for the CO2 compressors. This lowers the net HHV efficiency from 36.8% to 24.9%. This study assumed that new power plants would be Table 2.2 A comparison of the costs and efﬁciencies for a subcritical pulverized coal combustion power plant with and without carbon capture and sequestration (CCS).3 Without CCS With CCS

Net power output, MW Coal feed rate, t/hr Efficiency, HHV Plant cost, million $ Cost of electricity, cents/kw-hr Cost of CO2 emissions avoided, $/ton

550 219 36.8% 853 6.4

550 323 24.9% 1,591 11.9 68

23

24

Non-gasiﬁcation Uses of Coal

built, as opposed to modifying an existing power plant. This increases the coal feed rate for the CCS case was from 219 to 323 t/hr, a 47% increase, in order to maintain 550 MW of net output power. If a CCS system were added to an existing power plant, then the net power output would be reduced to about 68% of its former level. Widespread retrofitting of existing power plants would require substantial construction of new power plants to maintain power production levels. Adding a CCS system nearly doubles both the plant cost and the cost of electricity. Katzer et al.4 reviewed the CCS literature and concluded that adding CCS to coal-fired power production would about double the cost of electric power production. Distribution costs would not change, so CCS would increase residential power costs by about 50%. The enormous investment cost and steep electric power cost predicted for PCC with CCS prompted intense research into alternative CO2 absorbents and adsorbents. If a new power plant is to be built, then PCC with CCS is not the most economical approach to producing electricity with low CO2 emissions. The large number of existing power plants, however, provides a powerful incentive for adapting PCC technology. Alternative clean coal power production technologies are also being investigated. For example, much of the current interest in coal gasification is due to the predicted cost of electric power for coal-based integrated gasification combined cycle (IGCC) plant with CCS is substantially lower than PCC with CCS. Woods et al.3 studied IGCC with CCS using three different gasifiers. The concluded that electric power could be made for 10.3 cents/kw-hr. This is a substantial increase over PCC without CCS, but less than PCC with CCS. The high cost of clean coal technology has raised interest in power production technologies that produce less CO2.

OXY-COMBUSTION In simplistic terms, removal of CO2 from flue gas may be regarded as a CO2/N2 separation. The need for this separation may be eliminated if the furnace is fed oxygen instead of air. This is the basic concept of oxy-combustion. Figure 2.6 shows a simplified version of an oxy-combustion plant designed by Haslbeck et al.5 Oxygen is produced by an air separation unit (ASU). Typically, this is a cryogenic air distillation process; but other air separation techniques, such as pressure swing adsorption, have also been used. The oxygen purity from the cryogenic distillation unit is 95%. The impurities are argon, 3.4% and nitrogen, 1.6%. Flames fed nearly pure oxygen are much hotter than flames fed air. The materials of construction in the furnace cannot withstand these higher temperatures. Consequently, CO2-rich flue gas is recycled to the furnace to give an oxygen partial pressure that is comparable to air. Flue gas leaving the flue gas desulfurization unit is nearly saturated with water. So the flue gas is reheated slightly to avoid water droplets in the recycled flue gas.

Non-gasiﬁcation Uses of Coal

N2

Air

Air Separation Unit

Limestone, Water

O2 Coal

Pulverized Coal Combustion

Flue Gas Desulfurization

Ash

Gypsum

Flue Gas Recycle

Reheat Fan

Vent gas Cool Flash

Drier Compressor Water Pump Flue gas to sequestration

Figure 2.6 Oxy-combustion power plant based on the design by Haslbeck et al.5

A portion of the flue gas is withdrawn, compressed and then dried using a temperature swing adsorption unit. The flue gas is rich in CO2, but contains significant quantities of other gasses. This entire stream may be compressed and sequestered. Alternatively, the gas may be purified by cooling the gas and then separating liquid CO2 from a vent gas that contains most of the impurities. Table 2.3 shows the stream compositions when the purification process is used. The combined weight percent of the vent gas and the sequestered gas is less than 100% due to the small quantity of water removed from the moist flue gas. Note that the vent gas contains CO2 and SO2. These emissions can be eliminated if the purification process is removed. The need for CO2 purification depends on the gas specification limits required for sequestration.

25

26

Non-gasiﬁcation Uses of Coal

Fogash and White6 studied a process that would further purify the sequestered CO2 and reduce the release of pollutants in the vent gas. They used multistage flue gas compression combined with interstage cooling and condensation of water. Most of the NOx in the flue gas consists of NO. The conditions of the compression train favor the oxidation of NO to NO2. Given sufficient residence time, NO2 reacts with SO2 to form SO3 this in turn, reacts with water to form sulfuric acid (H2SO4). The NO2 also reacts with water to form nitric acid (HNO3), and mercury reacts with nitric acid to form mercuric nitrate. Consequently, the bulk of the SO2, NOx and mercury in the flue gas leaves with the condensed water. The values shown in Table 2.3 are data from Haslbeck et al. for a supercritical pulverized coal combustion unit fed Illinois No. 6 bituminous coal. Haslbeck et al. noted that the flue gas desulfurization unit could be eliminated and that SO2 could be co-sequestered with CO2. This would substantially reduce capital and operating costs. They kept the flue gas desulfurization unit in their design because, without it, the recycled flue gas would increase the SO2 concentration to 3.4 to 3.5 times as high as the same unit without a flue gas recycle. With a high sulfur coal like Illinois No. 6, this would cause corrosion problems in the boiler. Haslbeck et al. suggested eliminating the flue gas desulfurization unit when a low sulfur coal, such as Powder River Basin coal, is used. Table 2.4, also using data from Haslbeck et al., compares the cost and efficiency of a pulverized coal combustion plant with and without oxy-combustion. The addition of an oxy-combustion system substantially lowers the efficiency and increases the cost of electric power production. Comparing the values in Table 2.4 and Table 2.2, we see that, compared to carbon capture and sequestration using amine absorption, oxy-combustion is a more efficient and less costly means of capturing and sequestering CO2. For a greenfield plant, the cost of power from an oxy-combustion plant is comparable to an IGCC plant. Oxy-combustion, unlike IGCC, can be retrofitted to an existing pulverized coal combustion plant. Table 2.3 Flue gas compositions when CO2 is puriﬁed by liquefaction.5 Component, % Moist ﬂue gas Vent gas

Sequestered gas

Ar CO2 H 2O N2 O2 SO2 Wt.% of flue gas Temp., C Press., MPa

0 95.85 0.01 1.46 2.67 0.01 80.87 21 15.17

3.66 83.40 0.21 9.81 2.92 0.01 100 104 3.35

19.31 31.10 0 45.56 4.01 0.01 18.94 9 3.21

Non-gasiﬁcation Uses of Coal

Table 2.4 A comparison of the costs and efﬁciencies for a supercritical pulverized coal combustion power plant with and without oxycombustion.5 The oxy-combustion case does not use CO2 puriﬁcation. Without oxy-combustion With oxy-combustion

Net power output, MW Coal feed rate, t/hr Efficiency, HHV Plant cost, million $ Cost of electricity, cents/kw-hr Cost of CO2 emissions avoided, $/ton

550 185 39.4% 868 6.32

550 249 29.6% 1,263 10.07 43

SARGAS Figure 2.3 illustrates that part of the difficulty in separating CO2 from flue gas is due to its low partial pressure. As will be shown later, one of the attractive features of IGCC is that CO2 is separated from a high pressure gas stream, with a higher CO2 partial pressure. A similar approach is used by the Sargas process,7,8 in which CO2 is separated from flue gas from a pressurized combustion process. A Sargas demonstration plant was installed at the Va¨rtan combined heat and power plant in Stockholm, Sweden. This plant uses a pressurized fluidized bed combustor (ABB Carbon P200 PFBC cycle). As shown in Figure 2.7, air is fed to a compressor/turbine on a common shaft. Air is compressed to about 1.3 MPa, and fed to the pressurized fluidized bed combustor. Coal is fed to the combustor as a coal/water slurry. Limestone fed to the combustor reduces SO2 emissions. The low bed combustion temperature, typically 850e880 C, reduces NOx emissions. A hydro-cyclone is used to remove fly ash from the flue gas. The flue gas then is cooled in a heat exchanger, and fed to a Benfield process9e11 to separate CO2. This process is similar to the MEA process shown in Figure 2.5, with potassium carbonate (K2CO3) used instead of MEA. The Benfield process, as marketed by UOP,11 also uses a proprietary soluble catalyst and a corrosion inhibitor. An advantage of the Benfield process is that, unlike amines, the inorganic chemicals used in the process do not degrade in the presence of oxygen. Flue gas is initially contacted with water to remove residual dust, NO2, and partially remove SO2. The gas then contacts the K2CO3 solution, where CO2 absorbs and reacts to form potassium bicarbonate (KHCO3). The bicarbonate solution is heated in a stripping column to decompose the bicarbonate, releasing CO2 and regenerating K2CO3. The CO2-free flue gas is used to cool the flue gas fed to the Benfield process. This also reheats the flue gas, which is then fed to the turbine side of the compressor/turbine set. Warm, low pressure flue gas leaving the turbine is used to preheat boiler feed water before it is vented to the stack. The flue gas turbine produces about 20% of the power generated by the plant. The other 80% is generated by the steam turbine.

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Non-gasiﬁcation Uses of Coal

Stack

CO2

Benfield process Flue gas

Steam

Pressurized fluidized bed combustion

Coal/ limestone slurry

Fly ash

CO2-free flue gas

Boiler feed water

Air

Compressor

Turbine

Generator

Figure 2.7 The Sargas process combines a pressurized ﬂuidized bed combustion process with a postcombustion CO2 removal process, here shown as the Benﬁeld process.

COAL-TO-LIQUIDS The production of liquid fuels from coal is expected to be a major application of coal gasification. The following is a brief description of coal-to-liquid processes that do not rely on gasification.

ENCOAL Coking is the oldest form of processing coal, other than simply burning it. Coal is heated in the absence of oxygen to produce solid coke, liquid coal tar, and a flammable gas. Coal tar was a major source of liquid fuel and chemical feedstock until petroleum became abundant in the 1950s and 1960s. The ENCOAL process12 is a mild coking process that was demonstrated in a 1000 ton per day plant near Gillette, Wyoming, in the 1990s. An updated version of this process is marketed by ConvertCoal.13 A goal of this process is to upgrade PRB subbituminous coal to a solid fuel product called Process-Derived Fuel (PDF); which has an

Non-gasiﬁcation Uses of Coal

HHV value comparable to bituminous coals from the eastern USA. Skov et al.14 show that the heating value of a Powder River Basin sub-bituminous coal can be increased from about 19.6 MJ/kg (8,400 Btu/lb) to about 26.4 MJ/kg (11,200 Btu/lb). This increase is primarily due to the removal of water from the coal, which is about 30 wt.% of the feed coal (see Equation 1.1). Only about 60% of the volatiles were removed during pyrolysis, so the PDF has ignition characteristics that are similar to bituminous coal. Drying and pyrolysis also remove much of the sulfur and mercury from the coal. The process also yields Coal Derived Liquid, which has characteristics similar to a No. 6 petroleum heating oil. This is a heavy oil that is used in industrial boilers. The gas produced by the process was burned to provide heat for coking. Run-of-mine coal is crushed and then screened to 2 1/8 inch. The coal is then dried and pyrolyzed at 538 C (1,000 F). The gasses are cooled to condense the CDL product, and the cooling temperature is just high enough to prevent the condensation of water. There have been numerous attempts to increase the heating value of PRB by drying it or, as in the case of ENCOAL, by mild pyrolysis. A recurring problem with these upgrading processes is that dried PRB tends to reabsorb moisture. Another problem is that dried PRB is prone to low temperature oxidation, which can lead to spontaneous combustion. As can be seen in Figure 1.2, when dried coal is exposed to oxygen, the tendency to oxidize upon further exposure to oxygen is reduced. The ENCOAL plant sought to reduce the tendency of PDF to spontaneously ignite by exposing the material to oxygen in a controlled fashion. Pyrolyzed solids were quenched, and then sent to a vibrating fluidized bed. There they were exposed to a gas with a controlled concentration of oxygen. This treatment was found to be insufficient, so PDF was further deactivated by spreading PDF on the ground in 30 cm (12 in) thick layers. This exposed the PDF to oxygen at ambient conditions. Each ton of PRB coal yielded about 1 ton of PDF and 1 barrel of CDL. Skov et al.15 gave a detailed analysis of the CDL. The specific gravity of this liquid is 1.06 (2 API gravity), so 1 barrel of CDL per ton corresponds to a 9 wt.% yield of CDL. This yield will vary with feed coal and pyrolysis conditions. Pyrolysis processes are attractive because of their simplicity, but liquid yields are typically low. The CDL was sold as a No. 6 heating oil replacement. This is a low grade petroleum product with a shrinking market. Consumers of No. 6 heating oil are switching to natural gas or coal. Petroleum refiners have made considerable investments to convert heavy oils to lighter, and more valuable, transportation fuels. CDL can be upgraded to produce more valuable liquids. The primary difference between CDL and petroleum products is that the CDL contains high levels15 of oxygen (12.5 wt.%), nitrogen (0.9 wt.%), and sulfur (0.2 wt.%); which are collectively known as heteroatoms in the petroleum industry. In a typical petroleum liquid, the intermolecular

29

30

Non-gasiﬁcation Uses of Coal

forces between the hydrocarbon molecules are predominately weak dispersive, also known as van der Waals, forces. The hetero-atoms in coal tar introduce stronger polar, induced polar, and hydrogen bonding forces. Removal of the hetero-atoms by hydrogenation results in large reductions in boiling point, density, and viscosity. Only a few hydrogenation tests have been carried out with CDL.15

Direct Hydrogenation of Coal The goal of most coal-to-liquids processes is to convert coal to a liquid that resembles a petroleum product. Compared to petroleum, coal: • • • • •

is a solid has a higher molecular weight has a lower hydrogen/carbon ratio has higher concentrations of hetero-atoms: oxygen, sulfur and nitrogen contains ash.

Catalytic hydrogenation of coal can produce a petroleum-like product. Hydrogen is added to carbonecarbon double bonds and aromatic rings to produce carbonecarbon single bonds. Large molecules are split, or hydrocracked, to produce lower molecular weight compounds. Oxygen, sulfur, and nitrogen are removed as water, H2S, and NH3. Hydrocracking, plus the removal of polar hetero-atoms; converts solid coal into a liquid, with some byproduct gas. The ash can be filtered from the liquid. A thorough description of coal hydrogenation processes would fill several books. Included below is a brief overview of the hydrogenation process. The Bergius process was used in Germany during World War II to produce liquid fuels from coal. Several related processes were intensively developed during the 1970s and early 1980s, but development halted when the price of crude oil dropped during the mid 1980s. The Shenhua16 direct liquefaction plant in China started operations in 2009. Coal hydrogenation catalysts are closely related to petroleum hydrogenation catalysts. For petroleum, the catalyst is typically a sulfided bimetallic catalyst on a porous ceramic base; generally Co/Mo, Ni/Mo, or Ni/W on a silica, alumina, or silica-alumina base. Since coal hydrogenation is a less mature technology, a wider range of catalysts have been used. In a porous, (solid) heterogeneous catalyst, fluid reactants diffuse into the pores, adsorb onto the surface, react, desorb, and then diffuse out of the pores. The obvious problem with coal hydrogenation is that coal is a solid, and cannot directly access the catalytic sites inside the catalyst pellet. This problem is solved using two mechanisms. The mechanism first is that the coal is slurried in a hot, heavy, highly aromatic recycle liquid. This partially dissolves the coal, and the dissolved coal molecules can contact the catalytic surfaces. Compared to petroleum hydrogenation catalysts, coal catalysts have large pores to accommodate the large coal molecules.

Non-gasiﬁcation Uses of Coal

naphthalene 2 H2

catalyst pellet coal

tetrahydronaphthalene

Figure 2.8 The donor solvent mechanism in catalytic coal hydrogenation.

The second mechanism is the donor solvent process, shown in Figure 2.8. An aromatic compound in the liquid, shown here as naphthalene, is partially hydrogenated, saturating one of the aromatic rings. This hydrogenated molecule then reacts with the solid coal surface. Hydrogen from the saturated ring is transferred to the coal surface, which reforms the aromatic ring. Petroleum hydrogenation is typically done at 4 to 20 MPa with a large stoichiometric excess of hydrogen. A gas/liquid separator after the reactor recovers the catalyst addition

gas/liquid separator expanded catalyst level

oil recycle coal/oil slurry distributor plate hydrogen catalyst, ash withdrawl ebullating pump

Figure 2.9 Ebullating bed hydrogenation of coal.

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Non-gasiﬁcation Uses of Coal

unconverted hydrogen, and this gas is compressed and recycled. A trickle bed reactor is used. Oil and hydrogen are fed to the top of a packed catalyst bed, and oil trickles through the catalyst. This type of reactor will not work when the feedstock contains solids, as it does during coal hydrogenation. Instead, an ebullating bed reactor (also called a slurry bubble bed), shown in Figure 2.9, is used. The catalyst pellets are fluidized by hydrogen and recycle oil. Since the catalyst is rapidly poisoned during coal hydrogenation, fresh catalyst need to be continuously added through the top of the reactor. Spent catalyst is withdrawn as a catalyst/ash/unreacted coal/oil slurry through the Recycle H2 Recycle compressor Makeup H2 Ebullating bed hydrotreater

Coal

Recycle solvent

Naphtha

Slurry tank Preheat Recycle H2

Distillate fuel oil Vacuum distillation

Recycle compressor

Liquids

Flexicoker

Coke to gasification

Trickle bed hydrotreater Preheat

Figure 2.10 Exxon Donor Solvent process for direct hydrogenation of coal to produce liquid fuels.

Non-gasiﬁcation Uses of Coal

bottom of the reactor. Zhang17 measured the rate of deactivation of a Co/Mo on alumina catalyst in a laboratory reactor. Reaction rates declined to thermal, noncatalytic, levels after hydro-treating about 1,000 g of Powder River Basin coal per gram of catalyst. An examination of the spent catalyst showed that the pores were filled with coke, and a calcium carbonate shell coated the exterior of the catalyst pellets. Figure 2.10 shows the Exxon Donor Solvent process,18 one of several direct coal hydrogenation processes. Coal is slurried with a heavy recycle oil, mixed with hydrogen, preheated, and then fed to an ebullating bed hydrotreater. The oil product from the hydrotreater is distilled to yield naphtha (unfinished gasoline), distillate fuel oil (diesel and home heating oil), and a heavy recycle solvent. The vacuum distillation bottoms are coked to produce liquids and coke that can be fed to a gasifier to produce hydrogen. The recycle solvent is hydrotreated in a trickle bed hydrotreater before it is recycled to the slurry tank.

REFERENCES 1. www.epa.gov/mercury 2. www.epa.gov/ttn/chief/ap42/ch01/final/c01s01.pdf. 3. Woods MC, Capicotto PJ, Haslbeck JL, et al. Cost and performance baseline for fossil energy plants, volume 1: Bituminous coal and natural gas to electricity, DOE/NETL-2007/1281, <www.netl.doe. gov/energy-analyses/pubs/Bituminous%20Baseline_Final%20Report.pdf>; 2007. 4. Katzer J, Ansolabehere S, Beer J, et al. The future of coal, options for a carbon-constrained world, an interdisciplinary MIT study 2007. 5. Haslbeck JL, Black J, Kuehn N, Lewis E, Rutkowski MD, Woods M, et al. Pulverized coal oxycombustion power plants. DOE/NETLd2007/1291 <www.netl.doe.gov/energy-analyses/pubs/PC Oxyfuel Combustion Revised Report 2008.pdf>; 2008. 6. Fogash K White V. Oxycoal combustion: opportunities and challengesdpurification of oxyfuelderived CO2. Am Inst Chem Engrs National Meeting, Salt Lake City, UT 2007. 7. Hetland J, Christensen T. Assessment of a fully integrated SARGAS process operating on coal with near zero emissions. Appl Thermal Engr 2008;vol. 28:2030-2038. 8. Bryngelsson M, Westermark M. CO2 capture pilot test at a pressurized coal fired CHP plant. Energy Procedia, 2009;vol. 1:1403-1410. 9. Benson HE, Field JH, Jimeson RM. CO2 absorption employing hot potassium carbonate solutions. Chem Engr Prog 1954;vol. 50:356-364. 10. Benson HE, Field JH, Hayes WP. Improved process for CO2 Absorption uses hot carbonate solutions. Chem Engr Prog 1956;vol. 52:433-438. 11. UOP. Gas processing: Benfield process, <www.uop.com/objects/99 Benfield.pdf>. 12. The ENCOAL mild gasification project, a DOE assessment DOE/NETL-2002/1171, <www.netl. doe.gov/technologies/coalpower/cctc/cctdp/bibliography/demonstration/pdfs/encol/netl1171.pdf>; 2002. 13. ConvertCoal, Inc., <www.convertcoal.com>. 14. Skov ER, England DC, Henneforth JC, Rinker FG. Synthetic crude oil and coal-char production by mild-temperature pyrolysis processing of low-rank coals. Am Inst Chem Engr Spring Nat Meeting; April 22-26, 2007. 15. Skov ER, England DC, Rinker FG, Walty RJ. Coal-tar chemicals and synthetic crude oil production from low-rank coals using mild-temperature pyrolysis. Am Inst Chem Engr Spring Nat Meeting; April 22-26, 2007.

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16. China Shenhua, <www.csec.com/index.en>. 17. Zhang T. Development of a catalytic coal liquefaction microreactor and testing of novel supports for coal liquefaction catalysts. PhD. Dissertation, University of Wyoming; 1994. 18. Technology Status Report: Coal Liquefaction, Technology status report 010. United Kingdom: Dept. of Trade and Industry; 1999.

CHAPTER

3

Gasiﬁcation Fundamentals Contents Process Goals Devolatization Reactions with Oxygen Char Reactions Additional Gas Phase Reactions Slagging Balancing Coal, Oxygen and Water Feeds Air Versus Oxygen Feed Estimating Syngas Composition from Equilibrium Calculations Example 3.1: Estimating Gasiﬁcation Temperature and Pressure based on Gas Composition Reaction Rates Heuristic Design of Gasiﬁers Rate-Limiting Steps Effect of Pore Size on Mass Transfer Rates Thiele Analysis Transition from Surface reaction Limited Rates to Mass Transfer Inhibited rates in Gasiﬁers Reaction Rate with Changing Solid Mass Random Pore Model Shrinking Core Model Langmuir-Hinshelwood Equations for Surface Reaction Rates Example 3.2: Effect of Excess steam or Excess CO2 on Gasiﬁcation Rate Activation Energies Gasiﬁcation Catalysts Experimental Techniques used to Measure Gasiﬁcation Rates Fluidization Regimes Computational Fluid Dynamic Models In Summary References

35 36 38 38 39 39 39 41 41 44 46 46 46 48 49 51 53 53 55 56 58 59 59 62 64 67 68 69

PROCESS GOALS Gasification is an incomplete combustion of coal or another solid feedstock. The primary goals of gasification are: 1. Convert the entire non-ash fraction of the feed to gas; 2. Produce gasses that preserve, as much as possible, the heat of combustion value of the feedstock. Coal Gasification and Its Applications. ISBN B978-0-8155-2049-8.10003-8, doi:10.1016/B978-0-8155-2049-8.10003-8

Ó 2011 Elsevier Inc. All rights reserved.

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Gasiﬁcation Fundamentals

Combustion converts the feed to gas, but these gasses cannot be further burned to produce energy. Gasses from a gasifier can be burned to produce energy, or chemically converted to other products.

DEVOLATIZATION When coal enters a gasifier, the first step is to dry the coal. H2 Oð1Þ/H2 OðgÞ DH o vap ¼ 43:99 kJ=gmole

R-3.1

This phase transition is written using the same format that will be used to describe chemical reactions. The vaporization of water can have a substantial effect on the overall gasification thermodynamics, especially when a high moisture, low grade coal is gasified, or when the coal is fed to the reactor as a coal-water slurry. With bituminous coals and a dry feed gasifier, the impact of water vaporization is relatively small. The coal also pyrolyzes to produce coke and volatiles: Coal/CokeðsÞ þ VolatilesðgÞ DH o rxn ¼ positive; variable

R-3.2

Reaction R-3.2 is a simplistic means of describing a complex and variable set of reactions. At temperatures greater than about 320 oC, carbonecarbon bonds, or bonds between carbon and oxygen, nitrogen, or sulfur, in the backbone of organic compounds break. In the beginning of the reaction, unstable molecular fragments are formed. Those fragments may further pyrolyze, or react to form relatively stable compounds. When cooled to ambient temperature, the pyrolytic products may consist of gas, liquids in the form of coal tar or solids in the form of soot. Because of the complexity of coal, and the dependence of pyrolysis yields on processing conditions; an empirical approach to modeling reaction R-3.2 is often taken. Fletcher et al.,1 developed a Chemical Percolation model for Devolatization (CPD) that uses structural data from coal and 13C NMR analysis to predict pyrolysis yields. Figure 3.1 illustrates the concepts used to quantify coal structure. Coal consists of aromatic clusters, which can consist of a single aromatic ring or multiple fused aromatic rings. These clusters are connected by aliphatic bridges. Side chains are aliphatic groups that may or may not be bridges. If a side chain is connected to the same aromatic group in more than one place, it is known as a loop. The parent coal is assumed to have an infinite molecular weight. Pyrolysis begins when a bridge fractures, yielding a nonvolatile, finite molecular weight fragment known as a metaplast. Coal, when pyrolyzed, typically goes through a semi-liquid intermediate phase; which decomposes to solid char and volatiles. The liquid fraction roughly corresponds to a metaplast. Electron microscope images2 of rapidly heated char particles show that the char particle is often a hollow sphere, which is formed by an expanding gas bubble in a semi-liquid matrix.

Gasiﬁcation Fundamentals

CH2 OH R R side chain

O CH2

CH2

bridges CH2

R

CH2

COOH

CH2 CH2

loop

aromatic clusters

R

Figure 3.1 Coal structural information measured using yield using the CPD model.1

13

C NMR and then used to predict volatiles

The metaplast undergoes one of two competing reactions. The first reaction is to split into two smaller fragments. The second reaction is to form char plus two smaller fragments. Volatility of the product fragments is described using Raoult’s Law and a correlation between pure component vapor pressure and molecular weight. A key assumption in the CPD model is that the reaction rate constants are independent of coal type. Instead, differences in pyrolysis yields are due to structural differences, which can be measured using 13C NMR. The model has been successfully applied to a wide range of coals, which is remarkable considering the chemical variability and complexity of coal. Since 13C NMR data for coal are not widely available, correlations have been developed for determining the CPD structural parameters using more commonly obtained coal assay data.3 Much of the volatile fraction further reacts to form soot. The chemistry of soot formation is not as well known.4 If the feed points for coal and oxygen, or air, are closely spaced in the gasifier, then the volatiles may burn to form combustion and partial combustion products: Volatiles ðgÞ þ O2 /CO2 þ CO þ H2 O þ H2 þ H2 S þ SO2 DH o rxn ¼ negative; variable

R-3.3

In a moving bed gasifier pyrolysis occurs at relatively low temperatures, typically 700 to 800 oC. Coal and oxygen flow through the reactor in a countercurrent fashion, so

37

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Gasiﬁcation Fundamentals

nearly all of the oxygen is consumed before it reaches the pyrolysis section of the gasifier. Consequently, this type of gasifier produces substantial quantities of coal tar. Fluidized bed gasifiers, operate at about 1,000 oC, and the feed points for coal and oxygen are often close. These gasifiers produce little tar, but substantial quantities of methane. Higher molecular weight volatile products are not produced because they are not stable at this operating temperature. Entrained flow gasifiers, which typically operate at 1,400 to 1,500 oC, produce no coal tar and very little methane.

REACTIONS WITH OXYGEN Oxygen is fed to the gasifier, either as a nearly pure stream of oxygen, or as air. Oxygen reacts with the char to produce carbon monoxide or carbon dioxide: 2CðsÞ þ O2 ðgÞ/2COðgÞ DH o rxn ¼ 221:31 kJ=gmole

R-3.4

CðsÞ þ O2 ðgÞ/CO2 ðgÞ DH o rxn ¼ 393:98 kJ=gmole

R-3.5

Note that the heats of reaction correspond to the way the reaction is written. For example, Reaction R-3.4 produces two moles of carbon monoxide, and the heat of reaction is twice the heat of formation of carbon monoxide. On the other hand, Reaction R-3.5 produces one mole of carbon dioxide, so the heat of reaction is the same as the heat of formation for carbon dioxide. Oxygen may also react with gas phase species. In addition to reaction R-3.3, the following reactions are significant: 2COðgÞ þ O2 ðgÞ/2CO2 ðgÞ DH o rxn ¼ 566:65 kJ=gmole

R-3.6

2H2 ðgÞ þ O2 ðgÞ/2H2 OðgÞ DH o rxn ¼ 484:23 kJ=gmole

R-3.7

2CH4 ðgÞ þ O2 ðgÞ/2COðgÞ þ 4H2 ðgÞ DH o rxn ¼ 71:44 kJ=gmole

R-3.8

Different forms of these reactions can be written. For example, the reaction of char to produce carbon dioxide, R-3.5, may be considered to be a combination of the reaction of char to form carbon monoxide, R-3.4, and the subsequent combustion of carbon monoxide, R-3.6. A reaction for the complete combustion of methane can be written by combining reactions R-3.6, R-3.7, and R-3.8.

CHAR REACTIONS In addition to the combustion of char, the following reactions are also significant: CðsÞ þ H2 OðgÞ/COðgÞ þ H2 ðgÞ DH o rxn ¼ þ131:46 kJ=gmole

R-3.9

Gasiﬁcation Fundamentals

CðsÞ þ CO2 ðgÞ/2COðgÞ DH o rxn ¼ þ172:67 kJ=gmole CðsÞ þ 2H2 ðgÞ/CH4 ðgÞ DH o rxn ¼ 74:94 kJ=gmole

R-3.10 R-3.11

Reaction R-3.9 is also known as the steam gasification reaction. Reaction R-3.10 is the CO2 gasification reaction, also known as the Boudouard reaction. Reaction R-3.11 is the methanation reaction.

ADDITIONAL GAS PHASE REACTIONS The water gas shift reaction is often used in a catalytic reactor downstream of the gasifier to adjust carbon monoxide and hydrogen ratios. In the gasifier, this reaction occurs spontaneously because of the higher reaction temperature: COðgÞ þ H2 OðgÞ4CO2 ðgÞ þ H2 ðgÞ

DH o rxn ¼ 41:21 kJ=gmole R-3.12

The steam methane reforming reaction is used in catalytic reactors to produce syngas from natural gas. In the gasifier, this reaction may occur spontaneously because of the higher reaction temperature: CH4 ðgÞ þ H2 OðgÞ4COðgÞ þ 3H2 ðgÞ DH o rxn ¼ 206:2 kJ=gmole

R-3.13

SLAGGING As described in Chapter 1, coal ash is a complex and variable mixture of minerals that melt over a range of temperatures. Fluidized bed gasifiers typically operate near 1,000 oC, which is just under the ash softening temperature. This maximizes feed conversion while avoiding temperatures that would make the ash sticky. Slagging gasifiers, such as entrained flow gasifiers and the BGL moving bed gasifiers, operate at 1,400 to 1,500 oC so that the slag will be sufficiently fluid. Melting the ash to form a slag involves an endothermic heat of melting: AshðsÞ/SlagðlÞ DH o melt ¼ positive; variable

R-3.14

BALANCING COAL, OXYGEN AND WATER FEEDS The complete combustion reactions, R-3.5, R-3.6, and R-3.7, are very exothermic release a great deal of heat, but the product gasses have no further combustion value. The steam and CO2 gasification reactions (R-3.9 and R-3.10), are endothermic; which means that the product gasses have a greater heating value than the reactants.

39

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Gasiﬁcation Fundamentals

To maximize the heating value of the syngas, we want to drive the endothermic gasforming reactions as much as possible. These reactions do not occur spontaneously, so we rely on exothermic reactions with oxygen to raise the mixture temperature to the desired gasification temperature and to provide heat for the endothermic gas-forming reactions. The endothermic steam gasification, R-3.9, and CO2 gasification, R-3.10, reactions can be driven by an external source of heat rather than by feeding oxygen to the gasifier. Sources of heat that have been considered include solar heat, nuclear heat, and external char or gas combustion. With an external heat source, heat transfer is complicated by selecting conductive heat transfer materials that will withstand high gasification temperatures and corrosive atmospheres. The typical feeds to a gasifier are coal, oxygen, and water. Oxygen may be fed as a nearly pure oxygen stream from an air separation unit, or as air. Water enters the gasifier as coal moisture, coal slurry water, or as steam. The combined oxygen and water feeds should be sufficient to completely gasify the feed. If the oxygen feed is excessive, then the reaction becomes more like combustion than gasification, and low heating value gasses are produced. The temperature is controlled by varying the oxygen/water balance. The reactions with oxygen are all exothermic, so oxygen tends to increase the gasifier temperature. The steam gasification reaction, R-3.9, provides additional gas formation. Steam gasification is endothermic, so it tends to reduce the gasifier temperature. To optimize gasifier operation one must find the correct (O2 þ H2O)/coal ratio and the correct O2/H2O ratio. Carbon dioxide can be used in place of steam, but the CO2 gasification reaction, R-3.10, is slower than the steam gasification reaction, R-3.9. Most syngas applications also favor the higher H2/CO ratio produced by the steam gasification reaction. The oxygen/water/coal ratios depend on the gasifier configuration, the operating conditions, and the choice of coal. For example, the BGL gasifier, a moving bed gasifier; coal enters at the top, and slag is removed from the bottom. Oxygen and steam enters near the bottom; and gas flows up through the bed of coal and out near the top of the gasifier. In a recent study,5 a bituminous coal, Illinois No. 6, was gasified. Hot gasses from the bottom of the gasifier rise through the incoming coal, preheating the coal, driving off volatiles, and drying coal. Consequently, the exiting syngas has a relatively low temperature, 537 oC. This energy efficient counter current design also results in a relatively low 0.54 oxygen/dry coal mass ratio. The water/dry coal mass ratio was 0.40. Steam was 81% of the water feed, and the balance was moisture in the feed. In another study,6 the same coal, was gasified using a Shell gasifier. This is a high temperature, entrained flow design. Coal and oxygen enter near the top. Syngas and slag are removed near the bottom. Syngas exits the gasifier at 1,427 oC. The higher operation temperature results in a relatively high 0.83 oxygen/dry coal mass ratio. The water/dry coal ratio was 0.16. Steam was 67% of the water feed, and the balance was moisture in the feed.

Gasiﬁcation Fundamentals

Since O2 and H2O are both sources of oxygen for gasification; one can compare molar oxygen, O, not O2, to carbon ratios for both gasifiers. This ratio is 0.9669 for the BGL gasifier and 1.015 for the Shell gasifier. Despite the large difference in operating conditions, the oxygen feed requirement is nearly the same. The difference is in the source of the oxygen. The O2/H2O mass ratio is 1.336 for the BGL gasifier and 5.145 for the Shell gasifier. Due to the higher operating temperature for the Shell gasifier; more of the oxygen must come from O2, rather than H2O. In some gasifiers, the water feed is greatly in excess of the stoichiometric quantity required to gasify the feed. In this case, the oxygen feed must also be increased to raise the excess water to the gasification temperature. This is especially true if water is fed as liquid water, rather than steam, because heat is required to boil the water.

AIR VERSUS OXYGEN FEED The simplest and least expensive source of oxygen for gasification is compressed air. This, however, introduces a large quantity of nitrogen and a smaller quantity of argon into the gasifier. Moisture-free air is about 20.95% oxygen, 78.08% nitrogen, 0.93% argon, and about 0.40% other gasses. Nitrogen and argon, which are largely inert, dilute the reactive gasses. Additional oxygen is required to raise these inerts to the reaction temperature. Syngas from an oxygen-blown gasifier is traditionally referred to as medium Btu syngas. The heat of combustion for this gas is typically 400 to 500 Btu/scf, 14 to 18 MJ/ standard m3. Syngas from air-blown gasifiers is known as a low Btu syngas. Dilution of the syngas by nitrogen and argon typically lowers the heat of combustion to 100 to 200 Btu/scf, 3.5 to 7 MJ/standard m3. Natural gas and methane made from syngas are high Btu gases with a heat of combustion in the 900 to 1,000 Btu/scf range; 32 to 40 MJ/ standard m3. If the syngas is to be burned for heat or power generation, and if carbon capture and sequestration are not required; then the simplicity of an air blown gasifier is desirable. For other applications, the nitrogen and argon diluents can present processing difficulties. Most of the anticipated gasifier installations include an air separation unit, or ASU, typically cryogenic air distillation, to provide oxygen for gasification. The oxygen purity is typically 95%. The main impurity is argon, with smaller quantities of nitrogen. The ASU is usually a substantial fraction of the overall gasification complex capital cost. A portion of the syngas is often burned to generate electric power to operate the ASU.

ESTIMATING SYNGAS COMPOSITION FROM EQUILIBRIUM CALCULATIONS Reaction rates may be considered as a rate of approach to chemical equilibrium. If the gasifier residence time is large compared to the rate of reaction; then the products leaving

41

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Gasiﬁcation Fundamentals

the gasifier will be near chemical equilibrium. Since many gasifiers produce nearequilibrium products, thermodynamic models are often used to predict gasifier performance. The final equilibrium state does not depend on the path taken to achieve equilibrium. This allows one to calculate the final state using an unrealistic mechanism. In commercial chemical process simulators, such as Aspen Plus, the equilibrium state may be calculated by using a hypothetical reactor to decompose coal into its elements; and then mixing the elements with the gas feed and calculating the final equilibrium state. Higman and van der Burgt published, along with their book,7 a computer program that predicts gasifier performance based on an equilibrium model. Gasifier ashes and slags typically contain a small, but significant quantity of unconverted carbon. The equilibrium model for a gasifier can be improved by history matching, meaning that the experimentally measured residual carbon is specified in the model. In addition to the gasesolid reactions in the gasifer, there are also gas phase reactions that affect the relative concentrations of the product gases. One of these is the water gas shift reaction (R-3.12). The equilibrium constant may be written in terms of the partial pressures of the reactants and products: KWG ¼

PCO2 PH2 PCO PH2 O

Eqn. 3.1

Figure 3.2 shows the effect of temperature on the equilibrium constant. The values shown were calculated using a Peng-Robinson equation of state. In entrained flow, slagging gasifiers, and operating temperatures are relatively high, producing a high CO/ CO2 ratio. In lower temperature gasifiers the CO/CO2 ratio is lower.

Figure 3.2 Effect of temperature on the water gas shift equilibrium constant. A constant 2 MPa pressure was assumed.

Gasiﬁcation Fundamentals

The following empirical equation predicts the water gas shift equilibrium constant, KWG, to within þ 0.003 of the calculated values. In this equation, T is in Kelvin. lnðKWG Þ ¼ 1:8907½lnðT Þ2 30:084lnðTÞ þ 117:942

Eqn. 3.2

Note that the water gas shift reaction reacts two moles of gas to produce two moles of gas. Since there is no net change in the moles of gas, pressure has very little effect on the equilibrium constant. For example, at 1,000 oC, the equilibrium constant is 0.6056 at 0.1 MPa, and 0.6027 at 6 MPa. An equilibrium constant for the methane steam reforming reaction (R-3.13) may be written as: KMS ¼

PCO ðPH2 Þ3 PCH4 PH2 O

Eqn. 3.3

Note that two moles of gas react to form four moles of gas. Consequently, high pressure tends to push this reaction backwards. Figure 3.3 shows the effect of temperature and pressure on this equilibrium constant. High temperatures push the reaction forward. Low temperature, non-slagging gasifiers tend to produce substantial quantities of methane. Very little methane is found in the syngas produced by entrained flow, slagging gasifiers. Equations 3.4, 3.5, and 3.6 are an empirical fit to the data shown in Figure 3.3. The pressure is in MPa and the temperature is in Kelvin. KMS ¼ eb P a

Eqn. 3.4

a ¼ 7:1635ðT=1; 000Þ5 þ 53:378ðT=1; 000Þ4 157:83ðT=1; 000Þ3 þ 230:86ðT =1; 000Þ2 166:32ðT =1; 000Þ þ 44:849

Eqn. 3.5

b ¼ 9:5578ðT =1; 000Þ3 52:5ðT =1; 000Þ2 þ 105:19ðT =1; 000Þ 61:45 Eqn. 3.6

Figure 3.3 Effect of temperature and pressure on the methane steam reforming reaction equilibrium constant.

43

44

Gasiﬁcation Fundamentals

Temperature and pressure measurements of a gasifier can be difficult to obtain. One may estimate an effective gasification temperature and pressure by measuring the composition of the syngas, and then calculating the equilibrium conditions that would create that syngas. From the concentration of H2, H2O, CO, and CO2, one may calculate the water gas shift equilibrium constant using Equation 3.1. One may then estimate the gasification temperature using Figure 3.2 or Equation 3.2. An additional measurement of CH4 concentration allows the pressure to be estimated. Note that the definition of a gas partial pressure is the mole fraction of that gas, y, times the system pressure. For component i: Pi ¼ yi P

Eqn. 3.7

This definition may then be substituted into Equation 3.3 to give: KMS ¼

yCO ðyH2 Þ8 p2 yCH4 yH2 O

Eqn. 3.8

Setting Equation 3.8 equal to Equation 3.4 gives: yCO ðyH2 Þ3 p2 ¼ eb P a lnðPÞ yCH4 yH2 O 1 ½b lnðyCO Þ 3lnðyH2 Þ þ lnðyCH4 Þ þ lnðyH2 O Þ ¼ 2a

KMS ¼

Eqn. 3.9

The constants a and b may be found from the temperature and Equations 3.5 and 3.6.

Example 3.1: Estimating Gasiﬁcation Temperature and Pressure based on Gas Composition Powder River Basin coal was gasified in a Transport Gasifier using an oxygen-enriched air feed to yield syngas with the following composition:8

H2O CO H2 CO2 CH4 Ar N2

mole% 27.5 7.65 12.61 12.77 1.92 0.12 37.45

Gasiﬁcation Fundamentals

The temperature and pressure of the gasifier will be estimated based on the gas composition. First, the water gas shift equilibrium constant will be calculated using Eqn. 3.1: KWG ¼

PCO2 PH2 yCO2 yH2 ð0:1277Þð0:1261Þ ¼ 0:76544 ¼ ¼ PCO PH2 O yCO yH2 O ð0:0765Þð0:275Þ

From Figure 3.2, we can see that the temperature is about 925 oC or 1,198 K. This approximate solution is useful, as Eqn. 3.2 has multiple roots: lnðKWG Þ ¼ lnð0:76544Þ ¼ 1:8907½lnðTÞ2 30:084lnðT Þ þ 117:942 This was solved using Solver in Microsoft Excel to find that T equals 1,184 K or 911 oC. Alternatively, we could have solved Eqn. 3.2 using the quadratic equation, and selected the root that was closest to the estimate based on Figure 3.1. The pressure can then be estimated using Eqn. 3.9. First, we calculate the temperature-dependent a and b constants from Eqn. 3.5 and Eqn. 3.6: a ¼ 7:1635ðT=1; 000Þ5 þ 53:378ðT =1; 000Þ4 157:83ðT =1; 000Þ3 þ 230:86ðT =1; 000Þ2 166:32ðT =1; 000Þ þ 44:849 ¼ 2:1768 b ¼ 9:5578ðT =1; 000Þ3 52:5ðT =1; 000Þ2 þ 105:19ðT =1; 000Þ 61:45 ¼ 5:4115 Then, using Eqn. 3.9: 1 lnðPÞ ¼ ½b lnðyCO Þ 3lnðyH2 Þ þ lnðyCH4 Þ þ lnðyH2 O Þ ¼ 2:1428 2a P ¼ e2:1428 ¼ 8:52 MPa The reported average riser exit temperature was 907 oC; which is very close to the estimated 911 oC based on the gas composition. The riser exit temperature should be a little lower than the temperature near the base of the gasifier because of endothermic reactions in the riser. On the other hand, the estimated pressure of 8.52 MPa, does not compare well to the reported pressure of 1.303 MPa. The good temperature prediction shows that the water gas shift reaction, which is shown in R-3.9, is nearly at equilibrium in the reactor. The poor pressure prediction shows that methane is formed during gasification and that the steam methane reforming reaction, shown in R-3.10, is not at equilibrium. In the high temperature environment of an entrained flow gasifier, the methane steam reforming reaction is more likely to be near equilibrium. As shown in Figure 3.2, however, the equilibrium methane concentrations are low, and obtaining an accurate measure of methane content may be difficult.

45

46

Gasiﬁcation Fundamentals

REACTION RATES Heuristic Design of Gasiﬁers When a typical petrochemical reaction is commercialized, a series of laboratory experiments are run to produce an accurate reaction rate equation. The engineer then designs the reactor using this reaction rate equation; and sets initial reactor conditions. Gasifiers, on the other hand, are designed without the benefit of a reaction rate equation. The initial design is based on heuristics, which are designs that are based on prior gasifier experience. After the gasifier is built, the operating conditions are empirically adjusted to obtain desirable operating conditions. Due, in part, to the lack of a reaction rate equation, the development of a new commercial gasifier typically takes several decades. Small scale tests are often not representative of commercial operation, so large scale tests are run. These experiments are large, lengthy, and expensive. This approach to gasifier development is driven by the complexity of the gasifier. In fluidized bed and entrained flow gasifiers, there is a turbulent flow of gas and solids within the gasifier; and there can be large variations in gas composition and temperature within the gasifier. The coal particles have complex and variable chemical composition and structure. Particle composition and structure changes as the particles react.

Rate-Limiting Steps Despite the complexity of gasification, progress has been made describing gasification rates. Devolatization and reactions involving oxygen (R-2.1 to R-3.5) are fast. The slower reactions that determine the gasifier residence time are the steam gasification reaction (R-3.7) and the carbon dioxide reaction (R-3.8). The direct hydrogenation of char (R-3.9) is slower, so it has a relatively small impact on the required residence time. As shown in Figure 3.4, a sequence of steps is required for a gas-solid reaction: 1. The reactant gasses move to the solid surface due to fluid flow and diffusion. 2. The reactant gasses adsorb on the solid surface. 3. The reactant gasses surface diffuse from the adsorption site to the reaction site, if required by the reaction mechanism. 4. The adsorbed gasses and the solid react. 5. The product gasses surface diffuse from the reaction site to the desorption site, if required by the reaction mechanism. 6. The product gasses desorb. 7. The product gasses move into the bulk gas due to fluid flow and diffusion. Often, one of these steps is substantially slower than the other steps. This step becomes the rate-limiting step that controls the overall reaction rate. Any one of these steps may be the rate-limiting step. The remaining steps are nearly at equilibrium.

Gasiﬁcation Fundamentals

bulk gas

1

near surface

near surface

2 3

4 react

7

bulk gas

6 5

Figure 3.4 Potential rate-limiting steps when a gas reacts with a ﬂat solid.

log(reaction rate)

Figure 3.5, an Arrhenius plot, shows typical behavior for a gas reacting with a flat solid. At relatively low temperatures, on the right side of the plot, the surface reaction is the slowest step and controls the overall rate of reaction. As the reaction temperature increases; the surface reaction rate rapidly increases, until the surface reaction is no longer the slowest step. The rate of mass transfer then becomes the rate-limiting step. Compared to the surface reaction rate, the rate of mass transfer has a weak dependence on temperature; so the slope of the overall reaction rate with respect to temperature is less in the mass transfer limited zone. A low apparent activation energy is evidence for a mass transfer limited reaction. Steps 1 and 7 in Figure 3.4 can be broken down into several sub-steps. Figure 3.6 shows a conceptual break down for step 1. A breakdown for step 7 would be the same sub-steps in reverse order. The first sub-step consists of the movement of a reactant gas molecule from the bulk gas to the particle surface. This occurs by a combination of bulk gas flow and diffusion. But it is often modeled as if there is a stagnant boundary layer of gas around the particle, and the reacting gas molecule must diffuse from a well-mixed bulk gas phase across this layer to the particle surface. Most of the char reactive surface lies within the pores. It is reasonable to assume, for a range of conditions, that the reactive gas molecule will diffuse into the char pores rather than reacting on the char exterior particle surface. Since the gasification reactions

rate controlled by surface reaction rate rate controlled by mass transfer rate

1/T, K

Figure 3.5 Typical Arrhenius plot for a gas reacting with a ﬂat solid.

47

48

Gasiﬁcation Fundamentals

bulk gas

particle surface

Figure 3.6 Conceptual model of a gas reactant molecule traveling from the bulk gas, to the particle surface, through a macropore, mesopore, micropore, and ﬁnally adsorbing on a reaction site within the micropore.

produce more gas than they consume, there is a net flow of gas out of the char pores. The reactant gasses diffuse into the pores against this flow of gas.

Effect of Pore Size on Mass Transfer Rates Mass transfer rates in gasification are complicated by the pore structure of char. Coal and char contain macropores, with pore diameters greater than 50 nm, mesopores, with pore diameters between 2 and 50 nm, and micropores, with pore diameters between 0.4 and 2 nm. The rate of diffusion is affected by pore size. When the pore diameter is large compared to the mean free path of the gas molecules, then bulk diffusion is the dominate mechanism. This changes to slower Knudsen diffusion when the pore diameter is small compared to the mean free path of the gas molecules. The Knudsen diffusion constant, D, is given by: rﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃ 2 8RT Eqn. 3.10 D ¼ r 3 pM Where r ¼ pore radius R ¼ gas constant T ¼ temperature, K M ¼ molecular weight. Note that the diffusion constant decreases with decreasing pore diameter. The kinetic theory of gasses can be used to estimate the mean free path: h ¼ Where h ¼ mean free path, in cm x ¼ molecular diameter, cm P ¼ pressure, torr.

2:33 1020 T x2 p

Eqn. 3.11

Gasiﬁcation Fundamentals

Table 3.1 CO2 gas mean free paths estimated using kinetic theory. 1,000 oC

1,450 oC

0.1 MPa (1 atm) 6 MPa

509 nm 8 nm

376 nm 6.4 nm

This equation gives good results at low pressures (about 1 atm), but only approximate results at higher pressures where non-ideal gas behavior becomes significant. Table 3.1 shows estimated mean free paths for CO2 (0.322 nm diameter). The 1,000 oC temperature corresponds to a typical, non-slagging, fluidized bed gasifier; while the 1,450 oC temperatures corresponds to a typical slagging, entrained flow gasifier. Comparing these values to the pore diameters, we see that Knudsen diffusion is prominent in mesopores and some macropores at atmospheric pressure. At 6 MPa, near the upper end of the pressure range used in commercial gasifiers, bulk diffusion characterizes mass transfer in macropores and a mixture of bulk and Knudsen diffusion prevails in the mesopores. The Knudsen assumptions break down when the pore diameter is nearly the same size as the gas molecule diameter. A typical atomic diameter is 0.25 nm, so only small molecules may penetrate the micropores, and their movement is inhibited. Hurt et al.9 gasified chars prepared from a Utah sub-bituminous coal at 800 to 900 oC in carbon dioxide. They measured the mesoporous and microporous surface areas for the original char, and partially gasified chars. From these measurements, they concluded that most of the gasification occurred on the mesoporous surface, rather than the microporous surface. Slow diffusion rates for reactant and product gasses through the micropores probably prevented substantial reaction within the micropores.

Thiele Analysis Much of what we know about gasesolid reactions comes from the study of heterogeneous solid catalysts, which are typically porous solids that accelerate the reaction of fluid reactants to produce fluid products. The transition from a surface reaction rate limited reaction to a reaction that is inhibited by the rate of diffusion into pores is typically described by Thiele analysis. Note that this analysis does not consider the effect of pore size on diffusion rates. The effectiveness factor, h, is defined as: robs ¼ hrsl

Eqn. 3.12

Where robs ¼ the observed reaction rate rsl ¼ the surface limited reaction rate that would be observed with no mass transfer restriction. A value of h equal to one means that mass transfer rates do not affect the overall reaction rate. As the surface reaction rate increases, the rate of mass transfer restricts the overall reaction rate, and the effectiveness factor is less than one.

49

50

Gasiﬁcation Fundamentals

For a spherical particle, the effectiveness factor may be calculated from the following equation: 3 1 1 h ¼ Eqn. 3.13 f tanh ðfÞ f Where f ¼ Thiele modulus. For a first order reaction, the Thiele modulus may be calculated from: rﬃﬃﬃﬃﬃﬃ R kP f ¼ 3 D

Eqn. 3.14

Where R ¼ characteristic length ¼ sphere radius k ¼ first order reaction rate P ¼ reactant partial pressure D ¼ effective diffusion coefficient. The Thiele modulus may be thought of as a ratio of the surface reaction rate to rate of diffusion into the particle pores. Hodge et al.10 examined the impact of char morphology on reaction kinetics using an entrained flow reactor. In Eqns, 3.13 and 3.14, the char particle is assumed to be nearly spherical. Hodge et al. and others observed that the char particles frequently contain large bubbles. Coal, as it pyrolyzes, often goes through a semi-fluid state, and large bubbles can form in the particle before it completely solidifies. Hodge et al. classified their char particles using a system proposed earlier by Benfell et al.,11 which is illustrated in Figure 3.7. Group I particles are solid bubbles, in which a large, central void is encased by a porous carbon sphere. Group III particles contain no large bubbles. Group II is an intermediate conformation, which include both large bubbles and regions with small pores. Chars prepared from the three Australian coals tested by Hodge et al. contained 39 to 95% Group I particles, 0 to 61% Group III particles, and the balance consisted of 0 to 14% Group II particles. For Group III particles, the spherical assumption use for Eqns. 3.13 and 3.14 is reasonable. For Group I particles, the thin porous shell can be approximated as a flat slab. The Thiele modulus for a flat slab and a first order reaction is:

Group I porosity > 60 wall thickness < 5 µm

Group II porosity 40-60 wall thickness > 5 mm

Group III porosity < 40 wall thickness > 5 mm

Figure 3.7 Char morphology classiﬁcation according to Benfell et al.11

Gasiﬁcation Fundamentals

rﬃﬃﬃﬃﬃﬃ kP f ¼ L D

Eqn. 3.15

Where L is again a characteristic length, but this time this length is defined as the slab thickness. The usual definition of a Thiele modulus for a flat slab uses one half of the slab thickness as the characteristic length. This assumes that reactants diffuse into both sides of the slab. If the bubble for the Group I particle is intact, then reactants diffuse in to the bubble wall from only one side, the outside of the bubble, and the characteristic length is the bubble wall thickness. Note that this length tends to be much smaller than the characteristic length for a Group III particle, which is the particle radius. For a flat slab, the effectiveness factor is: h ¼

tanh ðfÞ f

Eqn. 3.16

The effectiveness factor for Group III particles is calculated using the spherical approximation, Eqn. 3.13 and Eqn. 3.14. The effectiveness factors for Group I and II particles are calculated using the flat slab approximation, Eqn. 3.15 and Eqn. 3.16. Group II particles aren’t readily approximated using an ideal shape, such as a flat slab, but the fraction of these particles appears to be small, so this approximation should not introduce large errors. The overall effectiveness factor is calculated12 from the effectiveness factor for each of the three char types: hoverall ¼ fI hI þ fII hII þ fIII hIII

Eqn. 3.17

hoverall ¼ Overall effectiveness factor

fI, fII, fIII ¼ Fractions of Group I, Group II, and Group II chars (see Figure 3.7)

hI, hII, hIII ¼ Effectiveness factors for Group I, Group II, and Group III chars.

Figure 3.8 is a plot of Eqn. 3.13 for spherical particles and Eqn. 3.16 for flat slabs. The surface reaction rate constant, k, is a strong function of temperature; while the diffusion coefficient, D, has a relatively weak temperature dependence. Consequently, increasing reaction temperature increases the Thiele modulus, moving to the right in Figure 3.8. For a Thiele modulus less than about 0.2, the effectiveness factor is nearly one, meaning that the overall reaction rate is surface reaction rate limited.

Transition from Surface reaction Limited Rates to Mass Transfer Inhibited rates in Gasiﬁers Kajitani et al.13 measured the carbon dioxide gasification rate for four coal chars, and found that mass transfer rates affected the overall reaction rate when the temperature exceeded 1,200 to 1,300 oC. The reduction in reaction rate due to

51

Gasiﬁcation Fundamentals 1 sphere

Effectiveness factor

52

flat plate

0.1 0.01

0.1

1

10

Thiele modulus

Figure 3.8 A plot of Eqn. 3.13 for spherical particles, and Eqn. 3.16 for ﬂat slabs. Increasing reaction temperatures increase the Thiele modulus, making it more likely that pore diffusion rates restrict the overall reaction rate (effectiveness factor less than one).

mass transfer restrictions was gradual, rather than the sharp bend shown in Figure 3.5. Below 1,200 to 1,300 oC, char particle sizes did not affect the gasification rate; but at higher temperatures, the gasification rates of smaller particles were faster than for larger particles. This is evidence for a transition from a surface reaction limited rate below 1,200 to 1,300 oC to a pore diffusion restricted rate at higher temperatures. Hodge et al.10 measure CO2 gasification rates at 1,000 to 1,400 oC. They found apparent activation energies of 78 to 156 kJ/mole, versus 242 to 281 kJ/mole for the same chars at lower temperatures. This is further evidence for diffusion restricted rates at the higher temperatures. Since non-slagging gasifiers typically operated at temperatures less than 1,100 oC; these gasifiers probably operate in the surface reaction rate limited regime. Slagging gasifiers, on the other hand, operate at 1,300 to 1,500 oC, so mass transfer rates probably affect gasification rates in these gasifiers. Low rank coals, which have relatively high reactivities, are better suited for nonslagging gasifiers than high rank coals. Tests with the Foster Wheeler Partial Gasification Module,14 a circulating fluidized bed gasifier, showed that 90% carbon conversion could be obtained with a subbituminous coal at 991 oC. With bituminous coals, the carbon conversion rate was 80% or less at temperatures as high as 1,077 oC. In the Transport Reactor,15 also a circulating fluidized bed gasifier, 88.5% carbon conversion was measured with a sub-bituminous coal at 860 oC, while carbon conversions were 69% or less with bituminous coals at temperatures up to 954 oC.

Gasiﬁcation Fundamentals

Entrained flow, slagging gasifiers, on the other hand, are know for their fuel flexibility. Feeds with low intrinsic gasification reactivity, such as petroleum coke, have been successfully gasified with this type of gasifier. If the overall rate of reaction is affected or dominated by mass transfer rates, then surface reaction rates are less important. Sharma et al.16 prepared ash-free coal extracts, called HyperCoals, by solventextracting three coal samples with 1-methylnaphthalene. Potassium carbonate, a gasification catalyst, was added to these extracts, and the mixtures were gasified in steam. The transition from a surface reaction limited rate to a mass transfer affected rate occurred at about 700 oC for extracts prepared from a sub-bituminous coal and a brown coal (lignite), and about 775 oC for an extract prepared from a bituminous coal. Steam gasification surface reaction rates are faster than carbon dioxide reaction rates. Further acceleration of the surface rate by a catalyst pushed the transition temperature to a relatively low level. Jess and Andresen17 found that the oxidation of chars is affected by mass transfer rates at temperatures greater than 500 oC.

Reaction Rate with Changing Solid Mass Unlike heterogeneous catalysis, the weight of the solid changes with time during gasification. When gasification rates are measured with a thermo-gravimetric analyzer (TGA), the rate of gasification is given by: 1 dw r ¼ Eqn. 3.18 w dt r ¼ reaction rate, fractional weight per time w ¼ sample weight t ¼ time. Since the sample weight changes with respect to time, the reaction rate is expressed as a rate of fractional weight loss.

Random Pore Model Not only does the weight of the sample change with respect to time, but the pore structure also changes with time. Bhatia and Perlmutter18 described this using a random pore model. As the gasification proceeds, the pores become larger. Initially, the pore surface area becomes larger. As gasification proceeds, pores merge and the walls dividing them disappear, which leads to a loss of surface area. The rate of reaction can be described in terms of the extent of reaction: pﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃ dx ¼ ð1 XÞ 1 jln ð1 XÞ Eqn. 3.19 dt

53

Gasiﬁcation Fundamentals

Where X is extent of reaction (0 initially, 1 when reaction is complete). s is dimensionless time, given by: s ¼

ks C n So t 1 30

Eqn. 3.20

Where ks ¼ surface reaction rate constant C ¼ concentration of reactant n ¼ reaction rate order So¼ initial surface area /unit volume, in units of cm2/cm3 (cm1). Note that ksCnSo is the initial reaction rate. 3o ¼ initial particle porosity j is a dimensionless structural property, given by; j ¼

4pLo ð1 30 Þ S02

Eqn. 3.21

Where Lo ¼ initial length of all of the pores in the particle, divided by the particle volume, typically about 3 106 cm/cm3. From a practical standpoint, we can measure So and eo, but not Lo. So j becomes a parameter used to fit experimental data, rather than a predictive parameter. Figure 3.9 shows a plot of Eqn. 3.19 with three values of j. The rate of the reaction is the slope of the curve. Initially, the rate of reaction increases because the pores become larger. Later, pore surface disappears and the rate of reaction slows. Liu et al.19 gasified Binxian coal in CO2 and measured the char surface area as a function of conversion. They found that the measured surface areas closely matched the trend predicted by the random pore model. The specific surface area (m2/g) reached a maximum at 35% conversion, and then declined. 1 0.9

Extent of Reaction

54

0.8

= 12

0.7 0.6

=4

0.5

=8

0.4 0.3 0.2 0.1 0 0

0.5

1

1.5

Dimensionless Time

Figure 3.9 Plot of Eqn. 3.19, the random pore model, showing how the extent of reaction increases with increasing dimensionless time. The effect of three different structural values, j, are shown.

Gasiﬁcation Fundamentals

The random pore model implies that surface reactivity is directly proportional to surface area, a common assumption for gas-solid reactions. Zhang et al.20 questioned that assumption, and suggested, that the surface reactivity is proportional to the number of active sites, which is not directly proportional to surface area. They measured the CO2 gasification rates of an Indonesian coal and an activated carbon using calcium and potassium catalysts. To measure the quantity of active sites, Zhang et al. measured the “irreversible” CO2 adsorption of a char. This is the quantity of CO2 per gram of char that absorbs onto the char at 300 oC and does not desorb when the ambient gas is changed to argon. The catalysts increased both the reactivity and the CO2 irreversible adsorption. With the potassium catalyst, the reactivity at high conversion rates was greater than predicted by the random pore model. This is an expected result, since the catalyst becomes more concentrated as the quantity of char decreases. With the calcium catalyst, on the other hand, the reactivity at high conversion was less than predicted by the random pore model. Zhang et al. attributed this to sintering of the calcium catalyst. They proposed a modification of the random pore model (Eqn. 3.19): pﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃ dx ¼ ð1 XÞ½ 1 jln ð1 XÞð1 þ qp Þ Eqn. 3.22 ds Where q ¼ cx if the reaction rate tends to increase, beyond that expected by the random pore model, as was the case for a potassium catalyst q ¼ (1cx) if the reaction tends to decrease faster than predicted by the random pore model, as was the case for a calcium catalyst. c, P are fitting parameters

Shrinking Core Model At very high reaction rates, the overall reaction rate is not limited by the surface reaction rate, or the rate of diffusion in the char pores. Instead, the overall rate of reaction is limited by the rate of mass transfer to and from the exterior of the char particle. In this case, all of the reactant gas is consumed either on the exterior of the coal particle or within the pores close to the particle exterior surface. This gives rise to the shrinking core model, an alternative to the random pore model, for describing how the rate of reaction changes with extent of reaction. In the random pore model, the particle diameter does not change. Instead, the porosity increases until nothing but ash is left. In the shrinking core model, the particle porosity is constant. The particle diameter decreases until the particle is gone. For the shrinking core model, the dependence of reaction rate on extent of reaction may be derived by assuming a spherical particle and noting that the mass is proportional to the particle volume and the reaction rate is proportional to the particle external surface area.

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Gasiﬁcation Fundamentals

4 3 pr a ð1 XÞ 3 dX rate a area ¼ 4pr 2 a dt dX a ð1 XÞ2=3 dt

mass a volume ¼

Eqn. 3.23

Like the random pore model, the shrinking core model predicts that the reaction rate falls as the particles are consumed. The shrinking core model does not predict the initial acceleration of the reaction rate, but there is often enough scatter in the data that either model can fit the data. While both models give similar predictions, the base assumptions are very different. The random pore model assumes a surface reaction rate limited mechanism; while the shrinking core model assumes a strongly mass transfer limited mechanism. In nonslagging gasifiers, the reaction is normally surface reaction limited, so the random pore model is preferred. High reaction rates, such as in a slagging gasifier or in combustion, are more likely to be mass transfer limited; which is the base assumption for the shrinking core model. To avoid confusion, much of the reaction rate data is reported at a common extent of reaction.

Langmuir-Hinshelwood Equations for Surface Reaction Rates Surface reaction rates, especially the effect of gas partial pressures on surface reaction rates, have been successfully modeled using Langmuir-Hinshelwood (LH) equations. The LH model does not include the effects of surface diffusion. For a more complete discussion of LH models, see a textbook on heterogeneous catalysis. Roberts and Harris21 used LH models to describe char gasification by carbon dioxide and steam (R-3.7 and R-3.6). Carbon dioxide adsorbs onto an active surface site, ), to produce gaseous carbon monoxide and a carbon-oxygen surface compound, C(O)(s). þCO2 ðgÞ/CðOÞðsÞ þ COðgÞ

R-3.15

The rate of this reaction is designated as r18. The reverse reaction rate, r19, is also significant. CðOÞðsÞ þ COðgÞ/ þCO2 ðgÞ R-3.16 The C(O)(s) surface complex may decompose to generate gaseous carbon monoxide and regenerate the active site at the rate r20. CðOÞðsÞ/ þCOðgÞ

R-3.17

If this fundamental mechanism is correct, then the overall carbon dioxide gasification rate is: rCO2 ¼

½Ct r18 PCO2 1 þ ðr19 =r20 ÞPCO þ ðr18 =r20 ÞPCO2

Where [Ct] is the concentration of active sites on the surface.

R-3.18

Gasiﬁcation Fundamentals

Chars made from three unspecified Australian bituminous coals were gasified in a thermo- gravimetric analyzer (TGA) at 900 oC and carbon dioxide pressures up to 3 MPa. The values of [Ct]r18 were 3.4 105 to 8.5 105 MPa1s1 and the values of (r18/r20) were 0.8 to 1.4 MPa1. The values of (r19/r20) were not determined because the concentrations of carbon monoxide in their experiments were insignificant. Since the reaction rates change as the char is consumed, all of these values were determined at a 10% conversion rate. Katajani et al.13 give a somewhat different LH reaction rate equation which includes a term for inhibition of the reaction by carbon monoxide. Roberts and Harris proposed a similar set of fundamental reactions for the steam gasification reaction. In the following reaction, it is not immediately apparent that the active site for steam adsorption, ), is the same active site shown for the carbon dioxide adsorption in R-3.15. þH2 OðgÞ/CðOÞðsÞ þ H2 ðgÞ

R-3.19

CðOÞðsÞ þ H2 ðgÞ/ þH2 OðgÞ

R-3.20

CðOÞðsÞ/ þCOðgÞ

R-3.21

The rates of reactions R-3.19, R-3.20, and R-3.21 are denoted as r22, r23 and r24, respectively. The rate of steam gasification then becomes: ½Cf r22 PH2 O rH2 O ¼ R-3.22 1 þ ðr22 =r24 ÞPH2 þ ðr22 =r24 ÞPH2 O Where [Cf ] is the concentration of active sites for the steam gasification reaction. For the three chars tested, values of [Cf ]r22 were 8.7 105 to 9.2 104 MPa1s1 and values of (r22/r24) were 0.7 to 4.5 MPa1. Values of (r23/r24) were not determined because the concentrations of hydrogen in their experiments were insignificant. Niksa et al.22 measured a strong hydrogen inhibition effect, but did not present enough data to calculate r23/r24. Temperature programmed desorption tests were run on the chars after partial gasification to investigate the surface oxide decomposition reactions (R-3.17 and R-3.20). The temperatures of the char samples were increased at a rate of about 0.17 oC/s, and the rate of carbon monoxide desorption was monitored as a function of time and temperature. The desorption peak for both the carbon dioxide gasification samples and the steam gasification samples occurred at nearly the same temperature, which suggests that the active sites for carbon dioxide gasification and steam gasification are the same sites. When both carbon dioxide and steam are present, the total gasification rate is not the sum of the carbon dioxide gasification rate (R-3.18) and the steam gasification rate (R-3.22). Instead, Roberts and Harris23 found that carbon dioxide inhibits the steam gasification rate. The combined gasification rate is: ðr18 =r20 ÞpCO2 R-3.23 rCO2 þH2 O ¼ rCO2 þ rH2 O 1 1 þ ðr18 =r20 ÞpCO2

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This reaction form is based on the assumption that carbon dioxide and steam compete for the same active sites. Adsorption of steam is blocked by pre-adsorbed carbon dioxide; so the rate of steam gasification is reduced by the presence of carbon dioxide. The rate of steam gasification is much faster than carbon dioxide gasification; so steam does not significantly retard the rate of carbon dioxide gasification. Again, potential inhibition of the gasification rates by carbon monoxide and hydrogen were not measured because there were insignificant quantities of these gasses in the experiments. Mu¨llen et al.24 developed a more complex set of LH rate equations to describe the gasification of a German bituminous coal.

Example 3.2: Effect of Excess steam or Excess CO2 on Gasiﬁcation Rate Coal is gasified at 900 oC and 70 bar in two different syngas streams, one containing excess steam, and the other containing excess CO2. The addition of steam or CO2 affects the concentration of the other gases via the water gas shift reaction, R-3.13. Mole% H2 O H2 CO CO2

Excess steam 44 20 13.2 22.8

Excess CO2 10 10 44.7 35.3

Reaction rates in these two atmospheres are calculated using R-3.18, R-3.22, and R3.13, using the following reaction rate parameters (these are the parameters for the most active char measured by Harris and Roberts21). The missing parameters are assumed to be zero. [Ct]r18 ¼ 0.12 bar1s1 r18/r20 ¼ 8.5 106 bar1 [Cf]r22 ¼ 0.47 bar1s1 r22/r24 ¼ 9.2 105 bar1 The calculated reaction rate in the excess steam environment is 16.3 s1 versus 6.25 s1 for the excess CO2 environment. Note that the excess steam rate is 2.6 times as fast as the excess CO2 rate. Fluidized bed gasifiers are sensitive to coal reactivity. Low grade coals tend to have higher gasification reactivities than higher grade coals, and this leads to higher carbon conversion rates in fluidized bed gasifiers. The reaction rate can be enhanced by adding excess steam, but this comes with a cost and efficiency penalty. First, there is the cost and energy required to generate steam. Next, additional oxygen is required to generate the heat needed to raise this excess steam to reaction temperatures. This shows that cost reductions and efficiency improvements are possible via better gasifier design or by the addition of gasification catalysts, either of which could reduce the need for excess steam.

Gasiﬁcation Fundamentals

Activation Energies The dependence of reaction rate on temperature is usually described using an Arrhenius equation: r ¼ k expðEa =RT Þ R-3.24 Ea ¼ activation energy R ¼ gas constant T ¼ absolute temperature R ¼ Pre-exponential constant The pre-exponential constant is sometimes called the frequency factor. Strictly speaking, every rate term in an LH rate equation should have its own activation energy and pre-exponential constant. The carbon dioxide gasification rate shown in R-3.15, should include three activation energies and three pre-exponential constants. Kajitani et al.13 report four activation energies and pre-exponential constants for the carbon dioxide gasification of each coal they tested. Mu¨llen et al.24 report twelve activation energies and twelve pre-exponential constants for their complete rate equation. In practice, however, the temperature dependence of gasification can be reported using a single activation energy. Kajitani et al. reported activation energies of 240 to 263 KJ/mole for the carbon dioxide gasification of the four chars that they tested. Roberts and Harris25 report 209 to 250 KJ/mole activation energy for carbon dioxide gasification of two chars and 221 to 235 KJ/mole for steam gasification of those same two chars. Ohtsuka and Asami26 found activation energies from 130 to 215 KJ/mol for steam gasification of eight chars.

Gasiﬁcation Catalysts Catalysts can be used to accelerate the surface reaction rate. Transition metal catalysts27 have been used, but alkaline metal salts are more commonly used. Kapteijn et al.28 measured the effect of group IA metals on the carbon dioxide gasification rate of Norit, an acid-washed, steam activated peat char. The reaction rate increased in the order of increasing metal atomic weight: uncatalyzed < Li < Na < K < Rb < Cs 29

They also measured the catalytic effect of group IIA metals, but did not see the same trend with respect to atomic weight: uncatalyzed < Be < Mg < SrzBa < CazK As shown in Figure 3.10, the catalysts had little impact on the activation energy, which is indicated by the slope of the lines, but a large impact on the pre-exponential constant. At 1,000 K (727 oC), the reaction rate for the highest calcium loading was about 600 times faster than the uncatalyzed Norit.

59

Gasiﬁcation Fundamentals 100

Norit

Ca

Li

Be

Rb

Ba

Rate, micromoles/mole carbon - s

60

Cs

Na

Mg

Sr

Ca Sr Ba

10 Li Mg

Be

K

Na K Rb

uncatalyzed Cs

1 0.75 1,060 o C

0.80

0.85

0.90

0.95

1.00 o

1.05 -1

1,000 X 1/T, K

1.10

1.15

1.20 560 o C

Figure 3.10 Arrhenius plot for catalytic carbon dioxide gasiﬁcation of Norit, an acid-washed peat char.28,29

Ohtsuka and Asami26 measured the catalytic effect of Ca(OH)2, at about a 5 wt.% Ca loading, on the rate of steam gasification of 16 coals with a range of coal ranks. The coals were gasified in 66 kPa steam at temperatures ranging from 600 to 700 oC. For each coal, there was a large increase in reaction rate when the catalyst was used. The catalyst had little effect on the activation energy, so the catalyst appears to increase the number of active sites on the coal surface. Levendis et al.30 measured the catalytic effect of calcium on the oxidation of a synthetic char. To be effective, the catalytic metal must be adsorbed on the char surface. The most common method of loading the catalyst is to soak the coal sample in an aqueous solution of the alkali metal salt, and then dry and coke the sample. Lang and Neavel31 found that they could enhance the calcium catalyzed steam gasification of Illinois No. 6 coal by preoxidizing the coal in air at 100 to 260 oC. This surface oxidation process created carboxylic acid groups, and the acidic hydrogen ion exchanged with calcium when the coal was mixed with a Ca(OH)2 solution. Another approach to achieving intimate contact between the catalyst and the char surface is to use a low melting point eutectic salt mixture as the catalyst. At gasification temperatures, the salt mixture melts and coats the char surface. This approach has been tested by McKee et al.32 and Sheth et al.33,34 Wang et al.35 simply mixed char with solid K2CO3 and steam gasified the mixture at 675 to 750 oC. Potassium is volatile at gasification temperatures, so potassium adsorbed onto the char surface from the vapor phase. Matsukata et al.36 classified gasification catalysts according to the behavior of the catalytic metal at gasification conditions. Group 1 catalysts, like potassium, vaporize.

Gasiﬁcation Fundamentals

Group 1 catalysts also diffuse into the char bulk. As gasification progresses, the quantity of catalyst on the char surface may decline due to vaporization and loss to the char bulk. Like group 1 catalysts, group 2 catalysts, such as barium, also diffuse into the char bulk. Group 2 catalysts differ in that they do not have a significant vapor pressure. Group 3 catalysts, such as calcium, do not have a significant vapor pressure nor do they diffuse into the bulk. Group 3 catalysts remain on the char surface throughout gasification. Non-volatile catalyst will leave the gasifier mixed with the ash. If the catalyst is watersoluble, then it may be recovered by leaching the ash. Some of the catalyst will be lost because the catalyst will react with the ash to form insoluble minerals; and because the leaching process is an imperfect separation. Catalyst separation from slag is not practical. Sharma et al.16 avoided the ash separation problem by gasifying ash-free Hypercoal extracts. Volatile catalysts will stay in the syngas until cools. Catalyst deposits may foul heat transfer surfaces. Since some catalyst will be lost, only inexpensive catalysts are economically viable. Cesium and rubidium salts, for example, are very effective catalysts, but too expensive. It’s unlikely that the eutectic salt mixtures tested by Sheth et al.33,34 will be economically feasible. There are a small number of catalysts that are sufficiently inexpensive such that recovery may not be necessary. Sodium carbonate (soda ash) and calcium hydroxide (hydrated lime) are good candidates for once-through catalyst use. Coals naturally contain catalytically active metals. Quyn et al.37 found that naturally occurring NaCl in Loy Yang coal, an Australian brown coal, strongly enhances gasification reactivity. Lang and Neavel31 acid washed a Wyoming sub-bituminous coal to remove the naturally occurring catalytic metals; and they found that the acid-washed coal had a much lower steam gasification rate. Few commercial gasifiers deliberately add a catalyst. In the early 1980s, the Exxon Catalytic Coal Gasification Process38 was under development. In this process, shown in Figure 3.11, coal, steam, and potassium carbonate are fed to an externally heated gasifier at 700 oC and 3.5 MPa. As can be seen in Figure 3.2, the equilibrium constant for the methane steam reforming reaction is low at these relatively low temperature, high pressure conditions, so the formation of methane is favored. The syngas is separated to produce a methane product stream, water, and a carbon dioxide waste stream. Hydrogen CO2

heat coal, catalyst, steam

separations

H2, CO, H2O

Figure 3.11 Exxon catalytic coal gasiﬁcation process.

water

CH4 product

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and carbon monoxide are recycled to the reactor. This is much different than the usual coal-to-methane process approach, in which methane is produced by reacting CO and H2 in a catalytic reactor downstream from the gasifier. Great Point Energy39 has a process called Bluegas, in which methane and carbon dioxide are directly produced in a coal gasifier using a proprietary catalyst.

Experimental Techniques used to Measure Gasiﬁcation Rates A typical gasification rate experiment starts by coking the sample. This is done, in part, because coal tars may interfere with the experimental procedure. For example, a small amount of coal tar can foul the column in a gas chromatograph used to analyze gas composition. A common experimental technique is the thermo-gravimetric analyzer (TGA). Char is loaded into a small pan in the TGA, and then the sample is heated, typically at about 10 oC/s, until the reaction temperature is reached. Reaction progress is monitored by continuously weighing the sample pan while gas with a controlled composition and pressure flows over the pan. To measure reactions R-3.15, R-3.19, and R-3.20, Roberts and Harris21,23,25 prepared a coal char by heating coal samples at 1,100 oC for three hours. They then ran TGA experiments over a range of temperatures, pressures, and CO2/H2O blends. A problem with this type of experiment is that chars are prone to sinter and lose surface area at coking conditions. Sintering will increase with increasing coking temperature and coking time. Megaritis et al.40 compared the carbon dioxide gasification reactivities of chars produced by heating a Daw Mill coal at 10 oC/s to 1,000 oC in a tubular reactor to chars prepared from the same coal, but heated at 1,000 oC/s to 1,000 oC in a wire mesh reactor. After 60 s at 1,000 oC in 20 bar CO2, about 40% of the rapidly coked sample gasified, versus only about 5% of the slowly coked sample. Roberts, Harris, and Hall41 compared the surface area and reactivity for chars made using their standard, slow, coking technique to chars rapidly coked in a pressurized entrained flow reactor at 1,100 oC at a range of nitrogen pressures. The rapid pyrolysis technique increased the mesoporous surface area (measured by nitrogen BET) by a factor of 5 to 8 times, and raised the microporous surface area (measured by carbon dioxide adsorption) by a factor of 20 to 50 times. Reactivities of the chars were compared by gasifying in a TGA at three conditions: 10 atm CO2 at 900 oC, 10 atm H2O at 850 oC, and 5 atm O2 at 350 oC. Carbon dioxide and steam gasification rates, on a mass basis, for the rapidly coked samples were about 30 times faster than the slowly coked chars. The authors also reported reaction rates based on microporous surface area, rather than weight, and found that they could explain most of the reaction rate differences by differences in surface area. However the surface area however, based rates still showed substantial differences due to the char preparation technique. The use of the microporous surface area, rather than the mesoporous surface area, conflicts with the findings of Hurt et al.9; who found that most of the gasification reaction occurs on the mesoporous surface area.

Gasiﬁcation Fundamentals

Lu et al.42 measured the effect of pyrolysis temperature on char physical structure and chemical composition for several Australian coals. The coals were rapidly pyrolyzed in a drop tube reactor at 900 oC, 1,200 oC, and 1,500 oC. X-ray diffraction measurements showed that the chars become more crystalline and ordered with increasing pyrolysis temperature. Surface areas of the chars were measured by nitrogen adsorption. The 900 oC and 1,200 oC chars had similar surface areas, but higher surface areas were measured for the 1,500 oC chars. With increasing pyrolysis temperature, H/C and O/C ratios decreased. Yu et al.43 and Yang et al.44 found that pyrolysis pressure strongly affects the char outside surface and macropore morphology. Large bubbles were visible in many of the char particles. Liu et al.45 measured steam gasification rates for a batch of coal injected into a preheated fluidized bed. Pyrolysis times, prior to gasification, were controlled by adjusting the time delay between coal injection and the start of steam flow. They found that coal reactivity declined with increasing pyrolysis time. Schurtz and Fletcher2 found that char particles produced by heating coal particles at 10,000 oC/s in a flat flame burner had different morphologies than char particles produced at 1,000 oC/s heating rates. At the higher heating rates, the char particles had mesoporous surface areas of about 300 m2/g, and did not tend to form the large bubbles observed at 1,000 oC/s. These results show that reaction rate data, gathered in the laboratory, reliably predict commercial gasifier rates only if the char is produced under conditions that approximate commercial gasifier conditions. Much of the existing reaction rate data, especially the older data, were gathered from chars that were slowly coked. This approximates the conditions in a commercial moving bed gasifier, but it is different than the conditions in a fluidized bed or entrained flow gasifier. For example, a small char particle in the Transport Reactor15, a fluidized bed gasifier, moving at the same velocity as the gas, has a single pass residence time of about 2 s. Similar residence times are found in entrained flow gasifiers. Rate data obtained on slowly coked chars are orders of magnitude slower than the reaction rates in fluidized bed and entrained flow gasifiers. Laboratory experimental techniques have been developed to measure gasification rates at high heating rates and short residence times. One technique is the drop tube furnace.46 Coal or char is fed to the top of a heated tubular furnace where it mixes with a preheated gas stream, and then falls to the bottom of the furnace. The gas and solids are separated at the bottom of the furnace, and the gas is analyzed to determine reaction rates. Kajitani et al.13 prepared a rapid-pyrolysis coke by heating coal in nitrogen at 1,400 oC in a drop tube furnace. This condition was an attempt to imitate the coal injection conditions in an entrained flow gasifier. Gasification rates in CO2 were then measured in a pressurized drop tube furnace at high temperatures, 1,200 to 1,400 oC, or in a pressurized thermo-gravimetric analyzer (TGA) at low temperatures, 800 to 1,000 oC.

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Results from both techniques fit on the same Arrhenius plot, showing that the slower TGA experiments did not undergo significant sintering over the lower temperature range. Another rapid heating technique is the wire mesh reactor. In this reactor, coal particles are held between two pieces of wire mesh. The sample is rapidly heated by passing an electric current through the mesh. Typical heating rates are 1,000 oC/s. The sample is then held at reaction conditions for 0 to 60 s. Power to the mesh is shut off, and the sample rapidly cools. The extent of coal conversion can be measured by comparing the initial and final weights of the sample. Messenbo¨ck et al.47,48 measured coal pyrolysis yields at 1,000 oC in flowing helium as a function of pressure and hold time. These experiments were then repeated with steam or carbon dioxide instead of helium. Weight loss in the reactive gasses was due to both devolatization and char gasification; so the reported gasification conversion was the weight loss in the reactive gas minus the weight loss at the same temperature and pressure in helium. These are some of the few recent laboratory gasification kinetic studies that were based on a coal sample, rather than a char sample. Zeng et al.49 extended the upper temperature limit in a wire mesh reactor from about 1,000 oC to about 1,500 oC to approximate the conditions in an entrained flow gasifier. The stainless steel mesh used by Messenbo¨ck et al. was replaced by a custom made molybdenum mesh. The mesh openings were less than 70 mm so that the char particles would be retained at high levels of conversion. Since molybdenum is much more conductive than stainless steel, a high current power supply was required. The flat flame burner technique used by Schurtz and Fletcher2 heats the injected coal particles at about 10,000 oC/s, a heating rate about one order of magnitude greater than a drop tube furnace or a wire mesh reactor. This approximates the conditions in the injector of an entrained flow gasifier. The heating rates in a drop tube or a wire mesh reactor better approximate the conditions in a fluidized bed gasifier; while the slow heating rates in a thermal gravimetric analyzer are a good approximation of the conditions in a moving bed gasifier.

FLUIDIZATION REGIMES So far, the discussion of gasification fundamentals has been limited to the particle level. Commercial gasifiers are large, complex reactors and the gasifier as a reaction system needs to be considered. Gasifiers are gas/solid contacting devices and are generally classified according to their fluidized bed characteristics. In a fluidized bed, shown in Figure 3.12, gas is injected into the bottom of a bed of particles. Gas flowing around the particle creates a drag force, pushing the particle upward; which is counteracted by gravity that pushes the particle downward.

Gasiﬁcation Fundamentals

gas particle bed

gas

distributor plate

gas

gravity drag force

Figure 3.12 In a ﬂuidized bed, gas ﬂowing through the particle bed pushes the particles upward, while gravity pushes the particles downward.

An overview of fluidized bed behavior is given by Grace50. If the upward drag force on the particles is less than the downward gravitational force on the particles, then the gas flow is not sufficient to suspend the particles. This is known as a fixed bed, and gas flows through the interstitial spaces in the bed. Moving bed gasifiers operate in the fixed bed regime. This sounds like a contradiction in terms. Solids are added to the top of the gasifier, and ash is removed from the bottom of the gasifier. As the solid feed is gasified, the bed slowly moves downward. Gas is fed to the bottom of the gasifier, and flows through the interstitial spaces in the particle bed. Consequently , although the bed is moving slowly, it operates in the fixed bed flow regime according to Grace’s fluidized bed classification. If the drag force exactly matches the gravitational force, then the gas flow is just sufficient to suspend the solids. This point is known as incipient fluidization. The bed expands. Particles at incipient fluidization tend to wobble, but they do not move quickly through the bed. Practical reactors do not operate at incipient fluidization because the exact balance between drag force and gravity does not allow a range of stable operating conditions. With increasing gas flow, assuming that other conditions remain the same, the reactor enters the bubbling fluidized bed regime. The continuous phase in the bed is a high density phase, much like incipiently fluidized bed. Rising through this bed are low density bubbles, which tend to stir the bed. A further increase in gas flow pushes the bed into the slugging regime. Very large bubbles, approximately the diameter of the reactor, are formed. This regime is generally avoided in commercial reactor operation. The turbulent regime occurs when the gas flow increases. The large bubbles of the slugging regime break up into unstable bubbles that rapidly change shape and size. In fast fluidization, the dense phase is no longer the continuous phase. The bed consists of streamers of high density and low density regions. Pneumatic conveying occurs when the gas velocity is high enough to sweep all of the particles out of the bed. There is no longer a dense phase. As mention earlier, moving bed gasifiers operate in the fixed bed flow regime. Entrained flow gasifiers operate in the pneumatic conveying flow regime. Some entrained flow gasifers flow top-to-bottom instead of the usual bottom-to-top direction assumed

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for fluidized bed beds. Fluidized bed gasifiers can be split into two types, bubbling fluidized beds; and circulating fluidized beds which operate in the turbulent or fast fluidization regimes. The fluidization regime can be predicted using two dimensionless parameters, U), a dimensionless gas velocity, and d)p, a dimensionless particle size. The dimensionless gas velocity is: 1=3 r2 ) U ¼ U Eqn. 3.24 ðDrÞgm U ¼ superficial gas velocity r ¼ gas density Dr ¼ rp e r rp ¼ particle density g ¼ acceleration due to gravity m ¼ gas viscosity. The dimensionless particle diameter, d)p, is given by; rðDrÞg 1=3 1=3 dp ¼ Ar ¼ dp m2

Eqn. 3.25

Ar ¼ Archimedes number dp ¼ particle diameter Once U) and d)p are obtained, one can predict the fluidization regime by examining the plot given by Grace50, or by the reproduction of this plot in standard handbooks.51 In fluidized bed gasifiers the coal is ground to very small sizes, typically around 100 microns. For such particles it can be shown that the incipient fluidization gas velocity is given by: ðDrÞgdp2 Eqn. 3.26 Umf ¼ 1650m Provided the particle Reynolds number Rep < 20, which is generally true for fine mesh coal particles. The coal particles will remain in one of the fluidization regimes until the gas velocity reaches the terminal settling velocity of the particles. Then they will be entrained and swept out with the flowing gas. The terminal settling velocity for fine mesh coal particles can be determined from: ðDrÞgdp2 Ut ¼ Eqn. 3.27 18m Provided the particle Reynolds number Rep < 0.4, which is also usually true or almost so for fine mesh coal particles. Thus the range of gas velocities that will keep the coal particles in a fluidized, non-entrained state is very large. The ratio of terminal settling

Gasiﬁcation Fundamentals

velocity, Ut, to the incipient fluidization gas velocity, Umf, is 1650/18 or about 91.6. Even for larger particles this ratio can be shown to be at least 8.72. For example, for coal particles of size of 100 microns fluidized in a gas of viscosity 3E-5 Pa$s, with a particleegas density difference of 2000 kg/m3 the incipient fluidization gas velocity would be approximately 0.004 m/s. The terminal settling velocity for these same conditions would be 0.36 m/s. A problem with this analysis is that the particle diameters and densities are not constant in coal gasification. The feed grinding equipment produces a range of particle sizes. As gasification proceeds, particle sizes and densities fall. This can lead to operational problems. For example, in a bubbling fluidized bed gasifier, small char particles are entrained in the gas flow, and swept out of the gasifier before gasification is complete. Consequently, carbon conversion levels are often less than desired.

COMPUTATIONAL FLUID DYNAMIC MODELS Gasifiers, as a chemical reaction system, are generally too complex to model using a simple kinetic expression. Instead, a computational fluid dynamic (CFD) model is generally needed to describe the complexities in a gasifier. In a simple CFD model, not a gasifier, the continuity equation and the equation of motion (see, for example, Bird, Stewart, and Lightfoot52) are discretized using a finite element method. In this process, the continuous equations are approximated by dividing the flow problem into a three-dimensional mesh. Each fluid mesh element is assumed to have uniform properties, and the mesh size is chosen such that the solution will approximate the continuously varying fluid properties. Equations are written to describe the geometry of the problem, and appropriate boundary conditions are chosen. This generates a large matrix of simultaneous equations. The solution of this matrix describes fluid flow in the system. Using a CFD model to describe a gasifier requires the addition of several layers of complexity. First of all, multiple phases are present. In addition to gas, the gasifier contains solids. The solid particles have a range of shapes and sizes, and the organic fraction of the solids is consumed as it passes through the gasifier. If a slagging gasifier is modeled, then liquids are also present. The phases in the reactor are complex and changing chemical mixtures. Both homogeneous and heterogeneous reactions must be considered. In addition to the continuity equation and the equation of motion, the energy equation and the mass transfer equation must be solved as coupled equations. Both of these latter equations include reaction terms. To accurately model the gasifier, the modeler needs to have complete knowledge of relevant phenomena; including solid properties, reaction rates, mass transfer rates, heat transfer rates, gas viscosities, and gas equation of state data. The CFD model is a complete catalogue of our knowledge, and lack of knowledge, about the system. Unfortunately,

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much of the necessary data is rarely available. In a gasification fundamentals workshop in 2009,53 it was estimated that we know about 90% of what we need to know to develop CFD models for combustion systems, and only about 20% of what we need to know to develop CFD models for gasifiers. Key areas requiring further research include heterogeneous reaction kinetics, the behavior of ash/slag, and interactions with the gasifier wall. CFD modelers typically use as much good, fundamental data as they have available, and then attempt to write reasonable models and approximations for the things they don’t know. For example, Bockelie et al.54 described a CFD model in which they used the kinetic models developed by Roberts, Harris, and co-workers despite the fact that the system they modeled used a different coal. A typical CFD model contains a large number of unknown parameters and questionable equations. The modeler will attempt to validate the model by comparing the model results to a limited set of experimental data. Since the model contains a large set of unknown, assumed parameters, it’s often possible to adjust those parameters to fit the experimental data. Even with a purely empirical model, using no fundamental information, it is possible to accurately match system performance over a limited set of conditions. The problem with an empirical model is that it cannot be reliably extrapolated to a new set of conditions. Empirical models also cannot be used to predict conditions at points within the system that cannot be readily measured. For example, a practical problem in entrained gasifier operation is the lifetime of the injector nozzle. The CFD model described by Bockelie et al. was used to estimate the effect of operating conditions on the temperature of the injector nozzle, which cannot be readily measured; using this model, the operating conditions were adjusted to give improved nozzle life. A CFD model based on a combination of known and unknown fundamental information should be regarded as a semi-empirical model. It is better than an empirical model, but not as good as a model based on well-known fundamental data. CFD models work well for combustion systems because the reactions in these systems are mass transfer limited. Even if we do a bad job of describing the heterogeneous reaction kinetics, we can still get a good model because the heterogeneous reaction kinetics have little impact on the overall reaction rates. Non-slagging gasifiers, on the other hand, appear to be kinetically limited. In these gasifiers, It is more important to understand the heterogeneous reaction kinetics than it is to understand mass transfer. Slagging gasifiers appear to operate in the coupled regime, where both heterogeneous kinetics and mass transfer strongly affect overall reaction rates.

IN SUMMARY Describing the fundamentals of gasification is somewhat like taking a picture with an out-of-focus camera. We’ve been tinkering with the camera for many years, and the

Gasiﬁcation Fundamentals

focus is better than it once was, but the picture is still a bit fuzzy. We are still tinkering with the camera, hoping for better pictures in the future. If we are successful, one day gasifiers will be designed and operated on the basis of fundamental knowledge. This should greatly reduce development time, and allow operation in a more efficient and cost-effective manner.

REFERENCES 1. Fletcher TH, Kerstein AR, Pugmire RJ, et al. Chemical percolation model for devolatization. 3. Direct use of 13C NMR data to predict effects of coal type. Energy & Fuels. 1992;6:414-431. 2. Schurtz R, Fletcher TH. Pyrolysis and gasification of a sub-bituminous coal at high heating rates, 26th Annual Int Pittsburgh Coal Conf, Sept. 20-23, 2009. 3. Genetti D, Fletcher TH, Pugmire RJ. Development and application of a correlation of 13C NMR chemical structural analyses of coal based on elemental composition and volatile matter content. Energy & Fuels. 1999;13:60-68. 4. Brown AL, Fletcher TH. Modeling soot derived from pulverized coal. Energy & Fuels. 1998;12:745757. 5. Bartone, Jr LM, White J. Industrial size gasification for syngas, substitute natural gas, and power production, DOE/NETL-401/040607, <www.netl.doe.gov/energy-analyses/pubs/Final%20Report %20DOE_NETL-401_1.zip; 2007>. 6. Woods MC, Capicotto PJ, Haslbeck JL, et al. Cost and performance baseline for fossil energy plants, volume 1: Bituminous coal and natural gas to electricity, DOE/NETL-2007/1281, <www.netl.doe. gov/energy-analyses/pubs/Bituminous%20Baseline_Final%20Report.pdf; 2007>. 7. Higman C, van der Burgt M. Gasification. 2nd ed. Elsevier; 2008. 8. Southern Company Services, Inc., Power Systems Development Facility Topical Report, Test Campaign TC 16, July 14eAugust 24, 2004, U.S.D.O.E. contract DE-FC21e90MC25140, <www.netl.doe.gov/ technologies/coalpower/gasification/projects/adv-gas/25140/TC16TopicalReport(2004).pdf> 9. Hurt RH, Sarofim AF, Longwell JP. The role of microporous surface area in the gasification of chars from a sub-bituminous coal. Fuel. 1991;70:1079-1082. 10. Hodge EM, Roberts DG, Harris DJ, et al. The significance of char morphology to the analysis of high temperature char-CO2 reaction rates. Energy & Fuels.; 2009. doi:10.1021/ef900503x. 11. Benfell K, Liu G-S, Roberts DG, et al. Proc Combust Inst. 2000;28:2233-2241. 12. Aris R. On shape factors for irregular particlesdI The steady state problem. Diffusion and reaction. Chem Eng Sci. 1957;6:262-268. 13. Kajitani S, Suzuki N, Ashizawa M, et al. CO2 gasification rate analysis of coal char in entrained flow coal gasifier. Fuel. 2006;85:163-169. 14. Robertson A. Development of Foster Wheeler’s Vision 21 partial gasification module. Morgantown, West Virginia: Vision 21 Program Review Meeting; Nov. 6-7, 2001. 15. Shadle LJ, Monazam ER, Swanson ML. Coal gasification in a transport reactor. Ind Eng Chem Res. 2001;40:2782-2792. 16. Sharma A, Saito I, Takanohashi T. Catalytic steam gasification reactivity of hypercoals produced from different rank of coals at 600e775 oC. Energy & Fuels. 2008;22:3561-3565. 17. Jess A, Andresen A-K. Influence of mass transfer on thermogravimetric analysis of combustion and gasification reactivity of coke. Fuel.; 2009. doi:10.1016/j.fuel 2009.09.002. 18. Bhatia SK, Perlmutter DD. Random pore model for fluid-solid reactions: I. Isothermal, kinetic control. AIChEJ 1980;26(no. 3):379-386. 19. Liu T-F, Fang Y-T, Wang Y. An experimental investigation into the gasification reactivity of chars prepared at high temperatures. Fuel. 2008;87:460-466. 20. Zhang Y, Hara S, Kajitani S, et al. Modeling of catalytic gasification kinetics of coal char and carbon. Fuel. 2010;89:152-157. 21. Roberts DG, Harris DJ. A kinetic analysis of coal char gasification reactions at high pressures. Energy & Fuels. 2006;20:2314-2320.

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22. Niksa S, Eckstrom D, Malhotra R. Chemical reaction kinetics for the initial stages of entrained-flow coal gasification, 26th Ann Int Pittsburgh Coal Conf, Sept. 20-23, 2009. 23. Roberts DG, Harris DJ. Char gasification in mixtures of CO2 and H2O: Competition and inhibition. Fuel. 2007;86:2672-2678. 24. Mu¨llen H-J, van Heek KH, Ju¨ntgen H. Kinetic studies of steam gasification of char in the presence of H2, CO2 and CO. Fuel. 1985;64:944-949. 25. Roberts DG, Harris DJ. Char gasification with O2, CO2, and H2O: Effects of pressure on intrinsic reaction kinetics. Energy & Fuels. 2000;14:483-489. 26. Ohtsuka Y, Asami K. Steam gasification of coals with calcium hydroxide. Energy & Fuels. 1995;9:10381042. 27. Holstein WL, Boudart M. Transition Metal and metal oxide catalyzed gasification of carbon by oxygen, water, and carbon dioxide. Fuel. 1983;62:162-165. 28. Kapteijin F, Abbel G, Moulijn JA. CO2 gasification of carbon catalyzed by alkali metals. Fuel. 1984;63:1036-1042. 29. Kapteijin F, Porre H, Moulijn JA. CO2 gasification of activated carbon catalyzed by earth alkaline elements. AIChEJ. 1986;32:691-695. 30. Levendis YA, Nam SW, Lowenberg M, et al. Catalysis of the combustion of synthetic char particles by various forms of calcium additives. Energy & Fuels. 1989;3:28-37. 31. Lang RJ, Neavel RC. Behaviour of calcium as a steam gasification catalyst. Fuel. 1982;61:620-626. 32. McKee DW, Spiro CL, Kosky PG, et al. Eutectic salt catalysts for graphite and coal char gasification. Fuel. 1985;64:805-809. 33. Sheth A, Yeboah YD, Godavarty A, et al. Catalytic gasification of coal using eutectic salts: Reaction kinetics with binary and ternary eutectic catalysts. Fuel. 2003;82:305-317. 34. Sheth AC, Sastry C, Yeboah YD, et al. Catalytic gasification of coal using eutectic salts: Reaction kinetics for hydrogasification using binary and ternary eutectic catalysts. Fuel. 2004;83:557-572. 35. Wang J, Jiang M, Yao Y, et al. Steam gasification of coal char catalyzed by K2CO3 for enhanced production of hydrogen without formation of methane. Fuel. 2009;88:1572-1579. 36. Matsukata M, Kikuchi E, Morita Y. A new classification of alkali and alkaline earth catalysts for gasification of carbon. Fuel. 1992;71:819-823. 37. Quyn DM, Wu H, Hayashi J-I, et al. Volatilisation and catalytic effects of alkali and alkaline earth metallic species during the pyrolysis and gasification of Victorian brown coal. Part IV. Catalytic effects of NaCl amd ion-exchangable Na in coal on char reactivity. Fuel. 2003;82:587-593. 38. Nahas NC. Exxon catalytic coal gasification process, fundamentals to flowsheets. Fuel. 1983;62:239241. 39. Great Point Energy, Our technology, ; 2008. 40. Megaritis A, Messenbo¨ck RC, Collot A-G, et al. Internal consistency of coal gasification reactivities determined in bench-scale reactors: effect of pyrolysis conditions on char reactivities under high pressure CO2. Fuel. 1998;77:1411-1420. 41. Roberts DG, Harris DJ, Wall TF. On the effects of high pressure and heating rate during coal pyrolysis on char gasification reactivity. Energy & Fuels. 2003;17:887-895. 42. Lu L, Sahajwalla V, Harris D. Characteristics of chars prepared from various pulverized coals at different temperatures using drop-tube furnace. Energy & Fuels. 2000;14:869-876. 43. Yu J, Harris D, Lucas J, et al. Effect of pressure on char formation during pyrolysis of pulverized Coal. Energy & Fuels. 2004;18:1346-1353. 44. Yang H, Chen H, Ju F, et al. Influence of pressure on coal pyrolysis and char gasification. Energy & Fuels. 2007;21:3165-3170. 45. Liu H, Zhu H, Kaneko M, et al. High temperature gasification reactivity with steam of coal chars derived under various pyrolysis conditions in a fluidized bed. Energy & Fuels.; 2009. doi:10.1021/ ef9004994. 46. Ouyang S, Yeasmin H, Mathews J. A pressurized drop-tube furnace for coal reactivity studies. Rev Sci Instrum. 1998;69:3036-3041. 47. Messenbo¨ck RC, Dugwell DR, Kandiyoti R. Coal gasification in CO2 and steam: Development of a Steam injection facility for high-pressure wire-mesh reactors. Energy & Fuels. 1999;13:122-129.

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48. Messenbo¨ck RC, Dugwell DR, Kandiyoti R. CO2 and steam-gasification in a high-pressure wiremesh reactor: The reactivity of Daw Mill coal and combustion reactivity of its chars. Fuel. 1999;78:781-793. 49. Zeng C, Chen L, Liu G, et al. Advances in the development of wire mesh reactor for coal gasification studies. Rev Sci Instrum. 2008;79. 084102. 50. Grace JR. Contacting modes and behaviour classification of gas-solid and other two-phase suspensions. Canadian J Chem Engr. 1986;64:353-363. 51. Green DW, Maloney JO, editors. Perry’s Chemical Engineers’ Handbook. 7th ed. McGraw Hill; 1997. 52. Bird RB, Stewart WE, Lightfoot EN. Transport Phenomena. John Wiley and Sons; 1960. 53. Richards G, Breault R. moderators, Gasification fundamentals workshop, 26th Ann Int Pittsburgh Coal Conf, Sept. 20-23, 2009. 54. Bockelie M, Denison, M, Swenson, D, et al. Modeling entrained flow gasifiers, 26th Ann Int Pittsburgh Coal Conf, Sept. 20-23, 2009.

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4

Gasiﬁers Contents Overview Moving Bed Gasiﬁers: The Lurgi Gasiﬁer BGL Gasiﬁer Fluidized Bed Gasiﬁers: The Winkler Gasiﬁer High Temperature Winkler Gasiﬁer U-Gas Gasiﬁer Foster-Wheeler Partial Gasiﬁer KBR Transport Gasiﬁer Entrained Flow Gasiﬁers: The GE Gasiﬁer ConocoPhillips E-Gas Gasiﬁer Shell Gasiﬁer Siemens Gasiﬁer Mitsubishi Heavy Industries (MHI) Gasiﬁer Pratt and Whitney Rocketdyne (PWR) Gasiﬁer Less Conventional Gasiﬁers: The Alter NRG Plasma Gasiﬁcation System References

73 73 75 77 79 79 79 80 83 88 90 91 92 93 96 98

OVERVIEW As discussed in the previous chapter, gasifiers are generally classified according to the fluidization regime in the gasifier; moving bed, fluidized bed, and entrained flow. Examples of each type of gasifier will be given in this chapter. Following this is a description of less conventional gasifiers. Commercial gasification has a long history, and a great number of gasifier designs have been developed. This chapter gives examples of gasifiers, but it is not a comprehensive description of all gasifier technologies.

MOVING BED GASIFIERS: THE LURGI GASIFIER The Lurgi gasifier is the oldest gasifier technology that is still widely used in commercial practice. In the early 1950s, Sasol, a South African firm, acquired the rights to use this gasifier from Lurgi, a German firm. Sasol still uses this gasifier to produce synthetic liquid fuels from coal. Because Sasol has made a number of improvements in the gasifier over the years, the Lurgi gasifier is often called the Sasol-Lurgi gasifier. In 2007, the Sasol-Lurgi gasifier was the most widely used gasifier in the world.1 The Sasol-Lurgi gasifier has been used with a full spectrum of coals, ranging from anthracite to lignite.2 Coal Gasification and Its Applications. ISBN B978-0-8155-2049-8.10004-X, doi:10.1016/B978-0-8155-2049-8.10004-X

Ó 2011 Elsevier Inc. All rights reserved.

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Figure 4.1 shows an illustration of the Sasol-Lurgi gasifier. The gasifier design solves the non-trivial problem of how to feed solid coal into a pressurized vessel, and how to remove solid ash from a pressurized vessel. Coal is fed to an atmospheric pressure bunker above the gasifier. When the bunker is full, a valve on the bottom opens, allowing the coal to drop into a coal lock. The valve then closes, and the coal lock is pressurized until it reaches the gasifier pressure, typically 2.4 to 3.5 MPa. The valve on the bottom of the coal lock opens, and the coal drops into the gasifier. Typically, each batch of coal weighs about 8 tons. Coal is held in the gasifier vessel for about an hour3 while oxygen and steam flow through the grate and into the coal bed. Figure 4.2 shows how the coal and gasses move counter-currently. This makes the Sasol-Lurgi design an especially energy efficient gasifier technology. The highest bed temperature, 615 to 760 C, occurs just above the grate, where the char is gasified. The hot syngas then rises and contacts cooler coal above the gasification zone. In the middle of the bed, rising hot gasses pyrolyze the coal, producing coal tar and char. In the top of the bed, the coolest region, coal is preheated and dried. Syngas and coal tar leave the gasifier at about 370 to 590 C. Compared to other gasification technologies, the operating temperatures in the Sasol-Lurgi gasifier are relatively low. Because of these low temperatures, a refractory lining is not required.

Figure 4.1 The Sasol-Lurgi gasiﬁer. Reprinted by permission from Sasol.

Gasiﬁers

coal

syngas + coal tar

pre-heating

pyrolysis

gasification (highest temperature)

grate

ash

O2, steam

Figure 4.2 Counter-current gas/solid ﬂow in the Sasol-Lurgi gasiﬁer. The bed is hottest at the bottom, and coolest at the top.

Because the bed must be free-flowing, only non-caking coals can be used. Fine coal cannot be used, because it will plug the interstitial spaces between the large coal particles. The feed coal is sized to about 3 to 30 mm.4 At the Great Plains Synfuels plant in Beulah, North Dakota, fine coal from the grinding circuit is sent to an adjacent pulverized coal plant. The Sasol-Lurgi gasifier produces a considerable quantity of tar, which complicates operations. Hot gasses leaving the reactor are quenched with a recycled water stream. The quench liquid is decanted to produce an organic liquid, which contains the bulk of the tar, and an aqueous layer. The aqueous layer contains water-soluble tar compounds, including phenol and cresylic acid (mixed isomers of methyl phenol). At the Great Plains Synfuels plant, tar is burned to produce steam for plant utilities, and the phenol and cresylic acids are recovered for sale. Newer, higher temperature gasification technologies often compare themselves to the Sasol-Lurgi gasifier and because of their higher operating temperature, little or no tar is produced. This is taken to be an advantage. Before crude oil became inexpensive in the 1950s and 1960s, coal tar was widely processed to make organic chemicals and liquid fuels. With increasing crude oil prices, coal tar may regain its former role as a liquid feedstock, and the tar yield from gasification may be regarded as an advantage, rather than as a disadvantage.

BGL GASIFIER The BGL (British Gas Lurgi) gasifier is a slagging version of the Lurgi gasifier. As shown in Figure 4.3, instead of a grate at the bottom of the gasifier, oxygen and steam are

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Figure 4.3 The BGL gasiﬁer. This is a slagging version of the Lurgi gasiﬁer, in which oxygen is injected into the slag layer.

injected into the slag layer using tuyeres. A conventional Lurgi coal feed system is used. Un-ground, as-received coal is fed to the gasifier. A stirrer near the top of the bed allows the use of caking coals. Because of the high temperature slag layer, greater than 99% carbon conversion is claimed. A refractory lining is required to withstand the high slag temperatures. In a process design by Bartone and White5, Illinois No. 6 bituminous coal is gasified in a BGL gasifier at about 2.7 MPa. Gas leaves the gasifier at 540 C, and is quenched by direct contact with water to 165 C. The condensed tar stream is about 8 wt.% of the coal feed. This tar stream is recycled back to the gasifier by injecting the tar, along with oxygen and steam, through the tuyeres and into the slag layer. Bartone and White also examined PRB subbituminous coal, and estimated about the same condensed tar yield. Table 4.1 shows the syngas composition estimated by Bartone and White, after quenching and further cooling. Note that the syngas contains 6 to 7 mole% hydrocarbons, primarily methane, which is the uncondensed tar fraction.

Gasiﬁers

Table 4.1 Estimated syngas compositions for the BGL gasiﬁer after a direct quench and further cooling.5 Coal component Illinois No. 6 mole% Powder River Basin mole%

Ar Benzene, toluene, xylene CH4 C 2 H4 C 2 H6 C 3 H6 C 3 H8 C 4 H8 C4H10 CO CO2 COS HCN H2 H 2O H2 S N2 Total hydrocarbons

0.0141 0.1033 5.9139 0.0471 0.2343 0.0190 0.0991 0.0147 0.0405 54.2938 4.2867 0.0619 0.0288 29.7464 0.2389 1.1521 3.7083 6.4719

0 0.1064 6.1009 0.1321 0.2415 0.0198 0.1019 0.0148 0.0396 55.9980 4.4245 0.0638 0.0297 30.7551 0.2390 1.1851 0.6398 6.7571

FLUIDIZED BED GASIFIERS: THE WINKLER GASIFIER The Winkler gasifier, commercialized in 1926, was the first industrial application of fluidized bed technology.6 The Winkler gasifier operates near atmospheric pressure, in the bubbling fluidized bed regime. Coal is ground to 0e8 mm. As shown in Figure 4.4, coal is fed to a bunker, and a screw feeder withdraws coal from the bunker and injects it into the bubbling fluidized bed. Since the bed is nearly at atmospheric pressure, a relatively simple coal feeding system can be used. The gas feed to the gasifier, consisting of steam and either air or oxygen, is split into two streams. Most of the gas is fed underneath the grate. This gas fluidizes and reacts with the solid bed. As the coal particles react, they become smaller and less dense. About 30% of the ash falls through the grate and is produced as bottom ash. The remaining 70% of the ash is entrained by the fluidizing gas and is carried into the head space. A major issue with this gasifier is that the entrained ash contains a significant quantity of unreacted carbon. The bubbling bed operates at about 1,000 C, a little below the ash softening temperature. The bed is operated near this upper operating temperature limit to maximize carbon conversion.

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syngas + fly ash hydrocyclone

steam coal bunker

water freeboard

screw feeder

fines return bubbling fluidized bed

coal steam + air or O2

rotating grate bottom ash

Figure 4.4 The Winkler gasiﬁer.

The remainder of the gas feed is injected into the freeboard. Oxygen reacting with the gas/solid mixture boosts the temperature to about 1,200 C, which partially melts the ash. The increase in temperature further increases the conversion of carbon in the entrained ash. Heat is removed from the top of the gasifier to re-solidify the ash before it leaves the gasifier. A hydrocyclone returns some of the entrained ash to the fluidized bed to further convert residual carbon in the ash. The Winkler gasifier was once widely used, but few, if any, commercial Winkler gasifiers continue to operate. The persistent issue of low carbon conversion appears to be reason for the demise of this gasifier technology.

filter bunker charge bin

gasifier

lock hoppers

compressor

Figure 4.5 Pressurized feed system for the High Temperature Winkler Gasiﬁer.8

Gasiﬁers

HIGH TEMPERATURE WINKLER GASIFIER The High Temperature Winkler gasifier was developed in the 1970s and 1980s. The primary change to the original design appears to be pressurized operation, at about 1 MPa.7 Pressurized operation increases gasification rates, which should improve carbon conversion. Pressurized operation required a change in the coal feed system, shown in Figure 4.5. Coal is loaded into an atmospheric pressure bunker. One of the two lock hoppers is then depressurized, and coal from the bunker falls into the lock hopper. The lock hopper inlet closes, and the lock hopper is pressurized. The bottom of the lock hopper opens, and coal falls into a line where it is pneumatically conveyed to the filter. Gas from the filter is recycled to the compressor, and coal falls into the charge bin. A solids metering valve on the bottom of the charge bin, which is pressurized, controls the flow of coal into the gasifier. The use of two lock hoppers allows pressurized coal to be continuously fed to the gasifier.

U-GAS GASIFIER The U-Gas gasifier, like the Winkler gasifier, is a bubbling fluidized bed gasifier. Significant differences include pressurized operation, and a conical screen instead of a rotating grate. The simplified flowsheets shown in the open literature do not show oxygen injection into the freeboard space. The gasifier operates at 0.3 to 3 MPa and 840 to 1,100 C. The U-Gas gasifier was developed by the Gas Technology Institute, and commercial licensing rights were acquired by Synthesis Energy Systems.9 At the time this book was written, three commercial U-Gas gasifiers were either operating (Hai Hua in Zaozhuang City) or under construction (Golden Concord in Inner Mongolia and YIMA in Henan Province) in China.10 A 100 ton/day bagasse, sugar cane waste, gasifier was built in Hawaii, and a 150 ton/day wood gasifier was built in Denmark11. Capital construction costs for the U-Gas gasifier are claimed to be lower than most competing gasifier technologies. This is an important consideration because the economic feasibility of most plants based on coal gasification is more strongly influenced by capital costs than by operating costs.

FOSTER-WHEELER PARTIAL GASIFIER The Foster-Wheeler Partial Gasifier,12,13 shown in Figure 4.6, is a fairly straightforward implementation of a circulating fluidized bed reactor. The goals for this gasifier were rather modest. Coal was to be partially gasified, and the remaining char would then be burned in a pulverized coal power plant. Carbon conversions between 45 and 80% were obtained for several bituminous coals. With sub-bituminous Powder River Basin coal, carbon conversions were 80 to 90% due to the higher reactivity of low grade coals.

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recycle cyclone

syngas cooler

syngas

gasifier body candle filter pre-cleaner cyclone

solids standpipe

char

char

coal, air, steam

Figure 4.6 The Foster-Wheeler Partial Gasiﬁer.

The gasifier body consists of a vertical pipe. Coal, air, and steam are injected into the bottom of the pipe, and the gas velocity is sufficient to carry the solids up. Near the top of the gasifier, the hot gas/solid mixture is sent to the recycle cyclone where coarse solids return to the bottom of the gasifier via a solids standpipe. The gasifier body, recycle cyclone, and solids standpipe are all refractory lined. Recycling the solids in this manner gives unreacted carbon in the coarse solids an additional opportunity to react. The inert fraction of the recycled solids has a thermal inertia role, dampening temperature variations in the gasifier. Gasses and fine solids leaving the recycle cyclone are cooled in the syngas cooler, and then sent to the pre-cleaner cyclone which removes a portion of the solid char. The remaining char is separated from the syngas in a candle filter. The gasifier operates at 0.7 to 0.9 MPa. Bituminous coals were gasified at 995 C to 1,065 C, and the subbituminous PRB coal was gasified at 945 to 995 C. A pilot scale gasifier, 12 m high by 18 cm I.D., was built and tested. No commercial units were built.

KBR TRANSPORT GASIFIER The Transport gasifier,14e19 shown in Figure 4.7, is a circulating fluidized bed gasifier. The main body of the gasifier has two sections, a larger-diameter mixing zone, on the bottom, and a smaller-diameter riser section, on the top. This differs from the FosterWheeler Partial Gasifier, shown in Figure 4.6, which has a constant diameter. The larger

Gasiﬁers

solids separation unit primary gas cooler riser

syngas

particulate control device

coal from lock hoppers air, O2 ,steam

mixing zone standpipe

Startup burner

ash depressurization

recycle syngas fly ash air, O2 ,steam ash depressurization

bottom ash

Figure 4.7 The Transport gasiﬁer.

diameter of the mixing zone lowers gas velocity, which allows more solids back-mixing and increases solids retention time. The gasifier is preheated with a gas-fired startup burner. The coal is dried sufficiently to remove moisture of the surface of the coal particles, so that the coal will flow readily through the feed system. Low grade coals often contain substantial moisture, much of it absorbed in the interior of coal particles, and this interior moisture does not need to be removed. The coal is ground to 250e600 mm average diameter,17 and fed to the gasifier using a pressurized lock hopper system. Coal is fed near the top of the mixing zone. Gas and reacting solids flow up through the riser. Coarse solids are removed by solids separation unit, and recycled via a standpipe with a J-leg seal on the bottom. Solids in the J-leg are fluidized by a recycle syngas stream.

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The original Transport gasifier design included two solids recycle loops. The first removed coarse solids using a dis-engager, and returned the solids via a standpipe and a J-leg seal. Medium solids were then removed with a cyclone and a short standpipe and loop seal that fed into the dis-engager standpipe. The design was later simplified to a single solids recycle loop. The solids separation unit in the new design appears to be a cyclone. Air or oxygen and steam enter the bottom of the reactor, just below where the recycle solids enter. The concept behind this configuration is that heat for the reactor would be primarily provided by burning carbon in the recycle solids. Analysis of the standpipes solids,14 however shows that the recycle solids generally contain less than 1 wt.% carbon. Temperatures within a gasifier tend to rise rapidly wherever oxygen is injected, because reactions with oxygen are fast and exothermic. As the flowing mixture moves away from the oxygen injection point, the reaction temperature tends to drop because of the endothermic steam and CO2 gasification reactions. A second air or O2 and steam injection point, just below the coal injection point, helps to level the temperature profile in the gasifier. Most gasification tests were air-blown. As shown in Table 4.2, syngas produced during O2-blown tests still contained large quantities of nitrogen due to the nitrogen used in the feed system and in purging instrumentation ports. Gas and fine solids leaving the solids separation unit pass through the primary gas cooler and into the particulate control device, which is a combined cyclone and candle filter. Gas enters the particulate control device tangentially, which throws some of the solids to the vessel wall, where they slide to the bottom. The remaining Table 4.2 Gas feed and product compositions for a Transport gasiﬁer fed PRB coal.14 Gas feed Air-blown, mole% Oxygen-enriched, mole%

Air O2 Steam N2 Riser exit temp, C Pressure, MPa Syngas Ar CH4 C2H6 CO CO2 H2 H2O N2

58.0 0 7.0 35.0 916 1.55

9.9 13.2 36.4 40.5 907 1.26

0.5 1.1 0.0 7.5 8.5 6.7 10.3 65.4

0.1 1.9 0 7.6 12.8 12.6 27.5 37.4

Gasiﬁers

solids are removed by banks of sintered metal tube filters in the center of the vessel. Solid filter cakes on the exterior of the tubes are periodically dislodged by backpressure pulses. Like other fluidized bed gasifiers, carbon conversion is generally higher for low grade coals than high grade coals. During air-blown operation, carbon conversion averaged 84% for an Illinois bituminous coal, 90% for Hiawatha bituminous coal, 95% for both PRB subbituminous coal and Freedom lignite, and 97% for Falkirk lignite18. Also like most fluidized bed gasifiers, much of the unconverted carbon can be attributed to fine particles that are quickly swept out of the gasifier. The bottom ash, which is essentially carbon-free, has a mean diameter of about 100 mm.17 The fly ash, on the other hand, has a mean diameter of about 10 mm. During gasification of PRB coal, the fly ash contained 20e40 wt.% unconverted carbon. Table 4.2 shows measured gas compositions from a process development unit14 for a Transport gasifier fed PRB coal operating in both air-blown and oxygen-enriched modes. Levels of C2H6 were not measurable, which shows that very little tar formed at these conditions. Methane, on the other hand, constituted a significant fraction of the combustible gas product. A considerable amount of nitrogen was fed to the gasifier through the coal feed system and by blanketing instrumentation taps. In a full-scale gasifier, the fraction of nitrogen fed to the gasifier would be reduced. In the oxygenenriched mode shown in Table 4.2, the O2/N2 ratio is closer to air than the air-blown mode, because of the effect of the N2 diluent. In the oxygen-enriched mode, a high steam feed rate was used. Due to the relatively low temperature, compared to entrained flow gasifiers, the water gas shift reaction favored conversion of CO and steam to H2 and CO2. The CO2/CO ratio was a relatively high 1.67, and the H2/CO ratio was a relatively high 1.65. A pilot scale Transport gasifier was built at the University of North Dakota. A larger process development facility was built and operated by the Southern Company in Wilsonville, Alabama. A previously-announced commercial IGCC plant19 based on a Transport gasifier was cancelled in 2009. This plant was to be an air-blown unit and fed PRB coal. No carbon capture and sequestration were included. The plant was cancelled due to uncertainties about greenhouse gas regulations. A commercial IGCC plant in Kemper County, Mississippi, using local lignite, will start operations in 2014. A 65% carbon capture and sequestration level is planned for this plant.

ENTRAINED FLOW GASIFIERS: THE GE GASIFIER The GE gasifier, shown in Figure 4.8, was originally developed by Texaco, an oil company that is now part of Chevron. This gasifier is known as the Texaco gasifier in older literature. The first commercial application was as an oil gasifier in 1956. The first commercial Texaco coal gasifier started operation in 1983.

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Figure 4.8 The GE (Texaco) gasiﬁer showing radiant heat recovery system.

The standard method of feeding coal to a high pressure GE gasifier is to finely grind the coal and then mix it with water to form a pumpable slurry. This slurry resembles the heavy oil fed to the first Texaco gasifier. The slurry and oxygen are injected into the top of the gasifier, and the gas/solid/slag mixture flows downward. There are two basic modes of operation. The first is the quench mode, illustrated in Figure 4.9. The hot gas/slag mixture is bubbled through a water bath, which solidifies the slag. The gas is cooled, and boiling water increases the steam content of the syngas. The slurry leaves the gasifier as a slag/water slurry. Note that solids are both fed into the pressurized gasifier, and are removed from the gasifier, as a water slurry. The gasifier is refractory lined. Hot slag slowly attacks the refractory liner, so the liner must be periodically replaced.20e22 This is a common problem in most slagging gasifiers. Figures 4.8 and 4.10 show the radiant heat recovery mode. A longer gasifier body is used, and steam tubes are embedded in the walls in the lower part of the gasifier body.

Gasiﬁers

O2

coal/water slurry

syngas quench water

slag/water slurry

Figure 4.9 The GE gasiﬁer in quench mode. coal/water slurry

O2

steam

steam

water

water syngas

quench water

slag/water slurry

Figure 4.10 The GE gasiﬁer in radiant heat recovery mode.

The estimated operating temperatures from Woods et al.23 are a 1,316 C syngas temperature, which is cooled to 593 C by the steam tubes, and further cooled to 210 C in the water quench. The estimated operating pressure is 5.6 MPa, which is a higher pressure than used by most other gasifiers. The radiant heat recovery mode has greater heat recovery, but the quench mode has a much lower capital cost. Frequently, gasification is followed by a water gas shift reactor

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to increase the H2/CO ratio. Steam must be fed to the water gas shift reactor, and the quench mode increases the steam content of the syngas. Table 4.3 shows syngas compositions estimated by Woods et al.23 for the radiant heat recovery version of the GE gasifier when fed Illinois No. 6 coal. Two stream compositions are shown. The “before quench” stream illustrates the effect of gasification conditions on syngas composition. The “after quench” stream shows the gas product from the gasifier, since the quench is an integral part of the gasifier. Argon constitutes about 0.9 mole% of air, and it is a common impurity in oxygen produced by cryogenic air distillation units, so nearly all syngas contains small quantities of argon. At the high gasification temperature, essentially no tar is produced and only a small quantity of methane is produced. Refer to Figure 3.3, which shows the effect of temperature and pressure on the methane steam reforming reaction equilibrium constant. The high gasification temperature also does not favor the conversion of CO and H2O to CO2 and H2 via the water gas shift reaction (see Figure 3.2), so the CO2/CO ratio is 0.44, a relatively low level. The quench, in addition to lowering the syngas temperature, also boosts the H2O syngas concentration from 14 to 27 mole%, which dilutes the other gas components. The coal/water slurry feed technique works well with bituminous coals, but with lower grade coals. The water/coal ratio is often far in excess of optimum due to the high intrinsic moisture in lower grade coals. To illustrate this point, the Wyoming State government commissioned a conceptual design of a 10,000 barrel/day PRB coal to liquids plant from a company called Rentech in 2005. They evaluated four different gasifiers, GE, Shell, ConocoPhillips and Future Energy (now Seimens), to incorporate into the design. Table 4.4 shows the results of the analysis for the four gasifiers. The GE gasifier required the highest coal feed rate, 13,600 ton/day for the GE gasifier compared to only 7650 ton/day for the Future Energy gasifier, because the higher water content of the slurry feed had to be vaporized by coal pyrolysis.

Table 4.3 Estimated syngas compositions23 for the GE gasiﬁer in radiant heat recovery mode when fed Illinois No. 6 coal. Component Before quench, mole% After quench, mole%

Ar CH4 CO CO2 COS H2 H 2O H2S N2 NH3

0.79 0.10 34.42 15.11 0.02 33.49 14.29 0.73 0.89 0.17

0.67 0.08 29.22 12.76 0.02 28.49 27.26 0.61 0.76 0.14

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Table 4.4 Screening results for a 10,000 barrel/day Fischer-Tropsch coal to liquids plant using Powder River basin sub-bituminous coal. Screening Results - Process Performance & Cost Once-Through FT Operation Producing 10,000 BPD Fuels GE ConocoPhillips Future Energy Shell (“Texaco”) (E-Gas) (GSP) (SCGP)

Coal Feed AR (STPD) Water in Coal AF (wgt.%) Oxygen Contained (STPD) Water Balance (GPM) Raw SG H2/CO ratio (molar) Net Power Export (MW) FT Products (BPD) EPC-CAPEX rating

13,600 55% 11,300 -185 1.29 64 10,000 4

9,150 48% 5,750 -200 0.88 100 10,000 2

7,650 10% 4,250 110 0.40 94 10,000 1

7,850 2% 4,450 200 0.42 81 10,000 3

To address this problem, GE acquired the Stamet pump technology.24,25 The Stamet pump, shown in Figure 4.11, is designed to pump dry solids against a pressure gradient. To do this, the pump uses a solids lock-up mechanism. If a large pressure gradient is applied to an unconfined bed of solids, the solids will fluidize and move towards lower pressure. If those same solids are pressed between two surfaces, then the solids are not free to fluidize, and the solids are in a lock-up state. The coal particles are compressed between pump surfaces and rotated from an atmospheric pressure hopper to a pressurized gasifier. Using this technology, the GE gasifier can be fed a dry coal feed obviating the need for the high water slurry fed operation previously used. GE expects that this technology will overcome its deficiency with respect to low rank, high moisture coals.

Figure 4.11 The Stamet pump is designed to pump dry solids against a pressure gradient.

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steam generator

syngas candle filter char

2nd stage coal/water slurry

1st stage coal/water slurry, O2

1st stage coal/water slurry, O2 quench water slag/water slurry

Figure 4.12 ConocoPhillips E-Gas gasiﬁer.

At the time this book was written, a pilot scale gasifier was under construction in Cheyenne, Wyoming.26 This pilot unit, jointly funded by GE and the State of Wyoming, is primarily intended to demonstrate the gasification of low grade coals using a GE gasifier equipped with a Stamet pump. In 2007, the GE gasifier was the second most commonly installed commercial gasifier1, the most common is the Sasol-Lurgi.

CONOCOPHILLIPS E-GAS GASIFIER The ConocoPhillips E-Gas gasifier was originally developed in the 1970s and 1980s by Dow Chemical.23 At that time, it was known as the Destec gasifier. As discussed in Chapter 3, a gasifier is normally fed a gas mixture consisting of steam and oxygen or air. Although rarely done, CO2 may substitute for all or part of the steam. The basic idea is to feed enough oxygen and steam to gasify all of the coal, and to adjust the oxygen rate to achieve the desired gasification temperature. An increase in the oxygen/steam ratio will increase the gasification temperature, but the syngas will have a lower heat of combustion. Entrained flow gasifiers operate at relatively high temperatures, which require relatively high oxygen/steam ratios. The basic concept of the E-gas gasifier is to use the high temperature syngas generated at slagging conditions to gasify coal injected at a second point. This more efficiently uses the heat generated by reactions with O2. As shown in Figure 4.12, a coal/water slurry and oxygen are injected into the high temperature first stage, operating at 1,316 to 1,427 C and 4.2 MPa.23 Slag formed at these conditions drops to

Gasiﬁers

the bottom of the gasifier, where it is quenched with water and withdrawn as a slag/water slurry. Hot syngas rises, passing through a restriction and into the second stage of the gasifier. A second coal/water slurry stream is injected at the inlet of the second stage; and the hot syngas provides heat for the endothermic pyrolysis, steam gasification, and CO2 gasification reactions. Gas and solids leave the second stage at about 1,010 C, which is below the ash softening temperature.23 Gasification of the second stage coal feed is incomplete, so the gas is cooled in a steam generator, the solid char is removed by a candle filter; and the char is returned to the first stage. The gasifier is refractory lined. The cold gas efficiency of the gasifier increases with increasing second stage slurry feed, and oxygen demand declines.27 At the Wabash River Coal Gasification Repowering Project,28 a commercial scale IGCC demonstration plant, the typical coal slurry feed split is 85% to the first stage and 15% to the second stage.27 Rutkowski et al.29 compared the cost of producing hydrogen from Pittsburgh No. 8 bituminous coal to PRB sub-bituminous coal. They chose an E-Gas gasifier, because this is one of the few gasifiers that has been demonstrated using both bituminous and sub-bituminous coals. A key assumption in their study was that both cases would have the same coal feed rate on a dry basis. They concluded that Pittsburgh No. 8 was the preferred feed, even though this coal is considerably more expensive than PRB. The cost of hydrogen was dominated by capital costs is the primary reason for the expense. In both cases, the capital cost was about the same. Pittsburgh No. 8, with its higher heating value, yielded more hydrogen product, which diluted the capital cost per unit of hydrogen product. Herbanek et al.,30 however, stated that lower grade coals, due to their higher reactivity, could be fed at higher rates to the E-Gas gasifier. They compared the costs of electricity from an IGCC plant and found that sub-bituminous PRB coal had a lower production cost than three bituminous coals. Table 4.5 shows syngas compositions estimated by Rutkowski et al. for Pittsburgh No. 8 and PRB coals. The relatively low temperature second stage boosts the methane concentration compared to the GE gasifier (see Table 4.3); although it is not obvious why the methane content differs for Pittsburgh No. 8 and PRB coals. The moisture content of the PRB coal, 26.6 wt.%, is much higher than the moisture content of the Pittsburgh No. 8 coal, 6.0%. This leads to higher water content in the PRB-derived syngas, which dilutes the other components. The higher water content in the PRBderived syngas and the relatively low temperature second stage also favors the conversion of CO and H2O to CO2 and H2 via the water gas shift reaction, so the PRB-derived syngas has a higher CO2/CO ratio than the Pittsburgh No. 8 derived syngas. The sulfur content of Pittsburgh No. 8 is much higher than PRB, 3.07 versus 0.82 wt.%, both dry basis, and this causes higher levels of H2S and COS in the syngas. In 2007, the ConocoPhillips E-Gas gasifier was the fourth most commonly installed gasifier.1

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Table 4.5 Estimated syngas compositions29 for the ConocoPhillips E-Gas gasiﬁer. Pittsburgh No. 8 Powder River Basin Coal component (bituminous), mole% (sub-bituminous), mole%

Ar CH4 CO CO2 COS H2 H2O H2S N2 NH3

0.82 0.42 41.95 9.75 0.04 33.20 12.19 0.78 0.57 0.28

0.75 0.17 26.61 16.04 0.01 28.34 27.17 0.19 0.52 0.21

SHELL GASIFIER Shell developed oil gasifiers in the 1950s. These units differ substantially from Shell’s coal gasifier, which was developed jointly23 with Krupp-Koppers from 1974 to 1981. After 1981, Krupp-Koppers offered a similar gasifier known as Prenflo.31 recycle compressor

quench gas gasifier

coal

fly ash

steam

steam syngas cooler

N2

grinding, drying, & feeding

boiler feed water boiler feed water

O2 slag

dry solids removal

fly ash (recycled)

Figure 4.13 The Shell coal gasiﬁer.

syngas

Gasiﬁers

The Shell coal gasifier23,32 is shown in Figure 4.13. The gasifier features a dry feed system. Coal is ground and dried, and N2 is used to pressurize the feed hoppers to the gasifier operating pressure, about 4.2 MPa. Coal is fed to opposite sides of the gasifier, near the bottom. Oxygen is fed to the gasifier below the coal injection points. Slag flows out the bottom. The hot gas/solid/slag mixture flows up through the gasifier at about 1,600 C. Near the top of the gasifier, a relatively cold recycle syngas stream, at about 200 C, is injected into the gasifier. This is shown as quench gas in Figure 4.13. The quench gas drops the syngas mixture to about 900 C. The steam boiler tubes in the syngas cooler cannot withstand the 1,600 C temperature of the unquenched syngas, so quenching drops the temperature to a tolerable level. The syngas cooler is taller than the gasifier. Shell also developed a water quench system to replace the syngas cooler.33 The fly ash contains unconverted carbon. This is recycled to the coal feed system. The Shell gasifier features a membrane wall, shown in Figure 4.14, rather than a refractory lining. The gasifier wall is cooled by boiling water to generate steam. A layer of solidified slag forms on the boiler tubes, and this solidified slag layer effectively functions like a refractory layer. Since the solid slag layer is the same composition as the liquid slag, the liquid slag does not attack the solid layer. Table 4.6 shows an estimated syngas composition for the Shell gasifier fed Illinois No. 6 bituminous coal. Because of the high gasification temperature and the lack of water fed to the gasifier, the syngas has very little methane and a high CO/CO2 ratio. In 2007, the Shell coal gasifier was the third most commonly installed gasifier, and the most popular gasifier for newly planned installations.1

SIEMENS GASIFIER Development of the Siemens gasifier started in the 1970s, and the first commercial plant was built in 1984. In 2006, Siemens acquired the gasifier technology from Future Energy.34,35 steam tube gasifier wall

atmosphere

solidified slag molten slag

hot syngas

Figure 4.14 Gasiﬁer membrane wall. Solidiﬁed slag replaces refractory lining.

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Table 4.6 Estimated syngas composition23 for the Shell gasiﬁer fed Illinois No. 6 coal. Component mole%

Ar CH4 CO CO2 COS H2 H2O H2S N2 NH3

0.97 0.04 57.16 2.11 0.07 29.01 3.64 0.81 5.85 0.33

Siemens, like GE, offers gasifiers, gas turbines, and steam turbines, all important components of an IGCC plant.36 The gasifier, shown in Figure 4.15, is also roughly similar to a GE gasifier in quench mode, shown in Figure 4.9. There are, however, significant differences. The Siemens gasifier features a dry feed system, using a lock hopper arraignment similar to that shown in Figure 4.5.37 The hoppers are pressurized with N2 or CO2 to the gasifier pressure, about 2.8 MPa.34,36 A gas fuel is used to preheat the gasifier. Coal and O2 are fed to the top of the gasifier, which operates at about 1,400 C. The gasifier is lined with a membrane wall instead of a refractory. Siemens also offers a refractory-lined version of their gasifier for low ash fuels. A wide variety of fuels have been gasified.38 Like the GE gasifier, the Siemens gasifier uses a water quench to solidify the slag, and the slag is removed as a slag/water slurry. The quench arraignment differs between the two gasifiers. In the GE gasifier, the syngas/slag mixture is blown into a water bath. In the Siemens gasifier, a water spray quenches the syngas/slag mixture.

MITSUBISHI HEAVY INDUSTRIES (MHI) GASIFIER Development of the MHI gasifier started in the 1980s.39 A 1,700 T/D demonstration plant in Nakoso, Japan started operations in 2007. A wide range of coals, including both bituminous and sub-bituminous coals, have been successfully gasified. The basic concept of the MHI gasifier is similar to the E-gas gasifier. The gasifier, shown in Figure 4.16, uses a split coal feed. Coal and air are fed to the bottom stage; and the rising, hot syngas is used to gasify a second coal feed stream. In the E-gas gasifier, these stages are known as the first and second stages. The equivalent nomenclatures in the MHI gasifier are the combustor and reductor stages.40 Endothermic reactions in the reductor section lower the temperature from the slagging condition to about 700 C, well below the ash softening temperature. Char is separated from syngas leaving the reductor stage and returned to the combustor stage.

Gasiﬁers

Figure 4.15 Siemens gasiﬁer.

The gasifier is air-blown. Most gasification plants that employ carbon capture and sequestration are oxygen-blown. Mitsubishi claims that the capital cost of their air-blown system is nearly the same as an oxygen-blown system, and that their efficiency is a little higher.41 The syngas contains about 30% CO and 10% H2. An oxygen-blown version of the MHI gasifier is being developed.42 The MHI gasifier uses a dry feed system and a membrane wall. Slag is removed from the bottom of the gasifier as a slag/water slurry. A very similar gasifier was developed by the Thermal Power Research Institute of China.43

PRATT AND WHITNEY ROCKETDYNE (PWR) GASIFIER The Pratt and Whitney Rocketdyne (PWR) compact gasifier, shown in Figure 4.17, was inspired by rocket engine technology. This gasifier was undergoing pilot scale development44 at the time this book was written. Commercial licenses are marketed by Zero Emissions Energy Plants (ZEEP).45

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syngas

reductor

char coal

combustor

air

slag water

slag/water slurry

Figure 4.16 Mitsubishi Heavy Industries (MHI) gasiﬁer.

Coal, O2

rapid mix injector

cooled membrane wall

rapid spray quench

syngas/slag

Figure 4.17 The Pratt and Whitney Rocketdyne Compact gasiﬁer.

Gasiﬁers

Pulverized coal and oxygen are split into multiple streams and fed to a rapid mix injector, which is then cooled. Tested injector pressures range from 1.5 to 6.7 MPa.46 Flame temperatures can reach 2,760 C, an extremely high temperature.47 The gasifier wall is cooled to keep the metal temperature below 540 C. Figure 4.18 shows a more complete PWR gasifier system. The feeding system features a dry solids pump,48 shown in Figure 4.19, similar in concept to the Stamet pump shown in Figure 4.11. The PWR solids pump features two opposed rotating belts. Coal particles are squeezed between the belts to achieve solids lock-up, allowing the solids to be pushed against a pressure gradient. Gas and solid slag leaving the gasifier first enter a hydrocyclone and then a bank of tube filters to remove the solid slag. A peer review49 of the PWR gasifier technology in 2006 identified the dry solids pump and the rapid mix injector as key items that require further technical development. Matuszewski et al.50 compared the PWR gasifier to the GE and Shell gasifiers and concluded that the PWR gasifier is both more efficient and less expensive. These results should be greeted with cautious skepticism, as Matuszewski et al. based their estimates on projected results from Pratt and Whiney for technology that had yet to be demonstrated. Still, such results encourage further development. Pratt and Whitney claims that their gasifier is only about 10% of the size of comparable commercial gasifiers, which should lessen capital costs. The high efficiency claims for the PWR gasifier are a little surprising when one considers the very high temperature and rapid gasification conditions. Normally, high low pressure hopper

quench

syngas hydrocylcone tube filter bank

dry solids pump gasifier

high pressure hopper

coarse slag

fine slag

Figure 4.18 Pratt and Whitney Rocketdyne (PWR) Compact Gasiﬁer System.

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Figure 4.19 The PWR dry solids feed pump.48 Coal is fed from a low pressure hopper above the pump, and discharged to a high pressure hopper below the pump.

efficiency processes require relatively low temperatures and a close approach to equilibrium, which lowers the driving forces, and leads to slow reactions. The high efficiency of the PWR gasifier may be due to the quench. The high temperature gasses produced by the upper part of the gasifier drive steam gasification reactions in the lower part of the gasifier; and the steam/solid slag stream leaves the gasifier at a relatively low temperature.

LESS CONVENTIONAL GASIFIERS: THE ALTER NRG PLASMA GASIFICATION SYSTEM The Alter NRG plasma gasifier51,52 uses a plasma torch to gasify a solid feedstock. This gasifier can be used to gasify coal, but it is especially attractive for difficult-to-gasify feedstocks, such as municipal solid waste (MSW). The largest current installation of the gasifier is the Eco-Valley Waste-to-Energy Facility in Utashinai, Japan, which gasifies 180 T/day of MSW and automotive shredder waste. At the time this book was written, a 120 MW coal power plant retrofit in Somerset, Massachusetts was planned,53 as well as a 40,000 BBL/day coal-to-liquids plant in Fox Creek, Alberta. Coskata54 built a semicommercial facility that gasifies wood chips using the Alter NRG plasma gasifier. The syngas is converted to ethanol in a bioreactor.

Gasiﬁers

Figure 4.20 Westinghouse Plasma Corporation torch used in the Alter NRG Plasma gasiﬁer.55

The heart of the process is a Westinghouse Plasma Corporation55 torch shown in Figure 4.20. Westinghouse Plasma Corporation is a subsidiary of Alter NRG. A plasma is an ionized gas and, in theory, a plasma can be made from any gas. Air appears to be the plasma feed gas for most gasifier applications.

Figure 4.21 Alter NRG Plasma gasiﬁer.51

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The plasma gasifier is shown in Figure 4.21. Solid feed enters through the side of the gasifier, and plasma torches are directed towards a moving bed of reacting solids. Near the plasma torch, temperatures can be as high as 3,000 C. Molten metal and slag leave the bottom of the reactor. Slagging operation is attractive for an MSW application, because toxic metals in the slag are essentially unleachable. Gases rise through the gasifier and enter an expanded freeboard zone, which allows coarser solids to fall back into the bed. The gas leaving the gasifier is about 900 to 1,000 C, which is hot enough to avoid tar. The heat of combustion of the syngas is about 80% of the feed.53 The syngas product can be burned to generate electricity. Electric power for the plasma torches is only about 2 to 5% of the feedstock energy input.

REFERENCES 1. Gasification World Database 2007: Current industry status. U.S. Dept. of Energy, Office of Fossil Energy, National Energy Technology Laboratory; 2007. 2. van de Venter E. Sasol-Lurgi coal gasification and low rank coal, Gasification Technologies Council 2005 Annual Conference. 3. Howard-Smith I, Werner GJ. Coal Conversion Technology. Noyes Data Corp; 1976. 4. Krichko AA. Theoretical basis of coal gasification. In: Oil and Gases from Coal. Pergamon Press for the United Nations; 1980. p. 89-124. 5. Bartone LM Jr, White J. Industrial size gasification for syngas, substitute natural gas, and power production, DOE/NETL-401/040607, <www.netl.doe.gov/energy-analyses/pubs/Final%20Report %20DOE_NETL-401_1.zip>; 2007. 6. Squires AM. Clean fuels from coal gasification. Science 1974;184:340-346. 7. Rezenbrink W, Wischnewski R, Engelhard J, Mittelsta¨dt A. High Temperature Winkler (HTW) coal gasification: A Fully-Developed for Methanol and Electricity Production, Gasification Technologies Council 1998 Annual Conference. 8. Bockelie M, Denison M, Swenson D, Senior C, Sarofim A. Modeling entrained flow gasifiers, 26th Ann Int Pittsburgh Coal Conf. Sept. 20e23, 2009. 9. Vail T. SES deployment & commercialization of U-GAS gasification technology Gasification Technologies Council 2007 Annual Conference. 10. Synthesis Energy Systems, Cleanly unlocking the value of coal: Our projects. <www.synthesisenergy. com/Projects.html>. 11. Lau F. Commercial development of the SES U-GAS gasification technology, Gasification Technologies 2009 Annual Conference. 12. Robertson A. Development of Foster Wheeler’s Vision 21 partial gasification module, Vision 21 Program Review Meeting, Morgantown, West Virginia, Nov. 6e7, 2001. 13. Engstro¨m F. Overview of power generation from biomass, Gasification Technologies Council 1999 Annual Conference. 14. Southern Company Services, Inc., Power Systems Development Facility Topical Report, Test Campaign TC 16, July 14eAugust 24, 2004, U.S. D.O.E. contract DE-FC21-90MC25140, <www.netl.doe.gov/ technologies/coalpower/gasification/projects/adv-gas/25140/TC16 Topical Report (2004).pdf>. 15. Shadle LJ, Monazam ER, Swanson ML. Coal gasification in a transport reactor. Ind Eng Chem Res. 2001;40:2782-2792. 16. Mann MD, Knutson RZ, Erjavec J, Jacobsen JP. Modeling reaction kinetics of steam gasification for a transport gasifier. Fuel. 2004;83:1,643-1,650. 17. Leonard R, Lambrecht RC, Vimalchand P, Yongue RA. Update on gasification testing at the power systems development facility, 32nd International Technical Conference on Coal Utilization & Fuel

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18. 19. 20. 21. 22. 23.

24. 25. 26. 27. 28. 29. 30. 31. 32. 33. 34. 35. 36. 37. 38.

Systems, June 10e15, 2007, <www.netl.doe.gov/technologies/coalpower/gasification/projects/advgas/25140/07ClGasP.pdf>. Southern Company Services, Inc. Power Systems Development Facility Topical Report, Test Campaign TC 17, October 26, 2004 e November 18, 2004, U.S. D.O.E. contract DE-FC21-90MC25140, <www.netl.doe. gov/technologies/coalpower/gasification/projects/adv-gas/25140/TC17 Topical Report (2004).pdf>. Pinkston T. Orlando Gasification Project: Demonstration of a 285 MW coal-based transport gasifier, Gasification Technologies Council 2006 Annual Conference. Johnson KI, Williford RE, Matyas J, Pilli SP, Sundaram SK, Korolev VN. Modeling slag penetration and refractory degradation using the finite element method, 25th Annual International Pittsburgh Coal Conference; 2008. Bennett JP, Kwong K-Y, Petty AV, Thomas H, Krabbe R. Interactions between slag and high chrome oxide refractory liners in air cooled slagging gasifiers, 25th Annual International Pittsburgh Coal Conference; 2008. Matyas J, Sundaram SK, Rodriguez CP, Edmunson AB, Arrogoni BM. Slag penetration into refractory lining of slagging coal gasifier, 25th Annual International Pittsburgh Coal Conference; 2008. Woods MC, Capicotto PJ, Haslbeck JL, Kuehn NJ, Matuszewski M, Pinkerton LL, et al. Cost and performance baseline for fossil energy plants, volume 1: Bituminous coal and natural gas to electricity, DOE/NETL-2007/1281, <www.netl.doe.gov/energy-analyses/pubs/Bituminous%20Baseline_Final %20Report.pdf>; 2007. Project Fact Sheet: Continuous pressure injection of solid fuels into advanced combustion system pressures. U.S. Dept. of Energy, Office of Fossil Energy, National Energy Technology Laboratory, <www.netl.doe. gov/publications/factsheets/project/Proj397.pdf>; 2006. Crew J. High efficiency feed system for PRB and other high-moisture feedstocks, Gasification Technologies 2008 Annual Conference. High Plains Gasification-Advanced Technology Center, <www.uwyo.edu/hpg-atc/>. Amick P. Conoco Phillips Technologies Solutions: Gasification update, Gasification Technologies Council 2004 Annual Conference. Wabash River Energy Ltd., Wabash River Coal Gasification Repowering Project: Final technical report, U.S. D.O.E. contract DE-FC21-92MC29310, <www.netl.doe.gov/technologies/coalpower/ cctc/cctdp/bibliography/demonstration/pdfs/wabsh/Final _Report.pdf>; 2000. Rutkowski MD, Buchanan TL, Klett MG, Schoff RL. Capital and operating cost of hydrogen from coal gasification, U.S. D.O.E. contract DE-AM26e99FT40465, <www.netl.doe.gov/technologies/ coalpower/gasification/pubs/pdf/HydrogenFromCoalGasificationFinalReport.pdf>; 2003. Herbanek R, Shah J, Gadde S, White J. Feedstock impact on an IGCC plant with CO2 capture, Gasification Technologies Council 2008 Annual Conference. Radtke K, Heinritz-Adrian M, Hooper M, Richards B. PRENFLO PSG and PDQ: Latest developments based on 10 years operating experience at Elcogas IGCC, Puertollano, Spain, Gasification Technologies Council 2008 Annual Conference. Zuideveld PL. Shell coal gasification process for power and hydrogen/chemicals, Gasification Technologies Council 2004 Annual Conference. Fournier G, Harteveld W, Prins M, Von Kossak T, Van den Berg R. A water quench for the Shell coal gasification process, 25th Annual International Pittsburgh Coal Conference; 2008. Volkman D. Future Energy Gmbh: An update on technology and projects, Gasification Technologies Council 2004 Annual Conference. Zwirn R. Gasification: Journey to commercialization, Gasification Technologies Council 2006 Annual Conference. Morehead H. Siemens IGCC and gasification activities: North America & China, Gasification Technologies 2009 Annual Conference. Klemmer K-D. The Siemens gasification process and its application to the Chinese Market, Gasification Technologies Council 2006 Annual Conference. Hannemann F, Schingnitz M, Lamp J, Wu B. Applications of Siemens fuel gasification technology for different types of coal, 25th Annual International Pittsburgh Coal Conference; 2008.

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39. Sakamoto K. Commercialization of IGCC/gasification technology for the US Market, Gasification Technologies Council 2008 Annual Conference. 40. Sakamoto K. Development of Mitsubishi air-blown IGCC technology with carbon capture, Gasification Technologies Council 2008 Annual Conference. 41. Fujii T. Deployment of IGCC technology with carbon capture, Sixth Annual Conference on Carbon Capture and Sequestration; 2007. 42. Sakamoto K. Mitsubishi IGCC project updates, Gasification Technologies Council 2009 Annual Conference. 43. Shisen X. GreenGendnear zero emission coal based power demonstration project in China, Gasification Technologies Council 2009 Annual Conference. 44. Darby A, Hartung J. Compact gasification system development status, Gasification Technologies Council 2009 Annual Conference. 45. Bernard B. Gasification for the next generation: Zero emissions energy plants, Gasification Technologies Council 2009 Annual Conference. 46. Hartung J, Darby A. Pratt and Whitney Rocketdyne (PWR) compact gasification system, Gasification Technologies Council 2007 Annual Conference. 47. Darby A. Status of the Pratt & Whitney Rocketdyne/DOE advanced single-stage gasifier development program, Gasification Technologies Council 2005 Annual Conference. 48. Sprouse KM, Matthews DR, Saunders T, Weber GF. PWR dry feed development status, 2008 International Pittsburgh Coal Conference. 49. Clayton S, Powell C, Rath L, Keairns D, Gray D, Geertsema A. et al. PWR gasifier peer review, <www.netl. doe.gov/technologies/coalpower/gasification/pubs/pdf/060221-a PWR PEER REVIEW REPORT_1. pdf>; 2006. 50. Matuszewski M, Rutkowski MD, Schoff RL. Comparison of Pratt and Whitney Rocketdyne IGCC and commercial IGCC performance, DOE/NETL-401/062006, <www.netl.doe.gov/technologies/ coalpower/gasification/pubs/pdf/060621 PWR Compact Gasifier Final Report_1.pdf>; 2006. 51. Alter NRG, <www.alternrg.ca>. 52. van Nierop P. Alter NRG Plasma gasification system for waste and biomass gasification, Gasification Technologies Council 2009 Annual Conference. 53. Bower R. Somerset plant refueling through plasma gasification, Gasification Technologies Council 2008 Annual Conference. 54. <www.coskata.com/facilities/>. 55. Westinghouse Plasma Corporation, A Division of Alter NRG, <www.westinghouse-plasma.com>.

CHAPTER

5

Underground Coal Gasiﬁcation Contents Underground Gasiﬁcation Concept Motivation Connections between Injection and Production Wells Process Control and Modeling Water Contamination UCG-recoverable Coal GasTech Process and Economic Study References

101 102 102 103 107 108 108 110

UNDERGROUND GASIFICATION CONCEPT The gasifiers described in Chapter 4 all consist of an industrial facility where a feedstock is brought to the gasifier. Underground coal gasification (UCG) takes a much different approach. In UCG, un-mined coal seams are reacted underground, with insufficient oxygen for complete combustion, to create syngas. Figure 5.1 shows a conceptual drawing of UCG. An oxidant, usually air, flows through an injection well and into a cavity in a coal seam. The oxygen and water within the coal seam react with the coal to produce syngas, which is withdrawn through a production well. Figure 5.1 shows the controlled, retractable injection point (CRIP) method, in which the injection point is gradually withdrawn as coal is consumed.

syngas

air surface

injection well

overburden production well

coal seam underburden

ash, rubble

Figure 5.1 Underground coal gasiﬁcation using a controlled, retractable injection point (CRIP). Coal Gasification and Its Applications. ISBN B978-0-8155-2049-8.10005-1, doi:10.1016/B978-0-8155-2049-8.10005-1

Ó 2011 Elsevier Inc. All rights reserved.

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MOTIVATION Much of the early commercial development of underground coal gasification1,2 occurred in the Soviet Union in the 1930s. The motivation for this development was ideological, rather than economic. Underground coal mining at that time was dirty, dangerous work and many miners were sickened, injured or killed by their work. Modern health and safety regulations in Western nations have greatly improved working conditions in underground mines, but in many developing nations, poor conditions remain. The current motivation for underground coal gasification is economic. Most known coal resources are too deeply buried to be economically mined in the near future (see the discussion of reserves versus resources in Chapter 1). The costs of drilling air injection and syngas production wells are a small fraction of the cost of a gasification complex, so deeply buried coal seams may be gasified at a reasonable cost. This has the potential of reclassifying a large quantity of known coal resources to economically attractive reserves. Currently, the cost of products based on coal gasification tends to be dominated by capital construction costs, rather than the cost of coal. Underground coal gasification avoids the cost of the gasifier, which is usually a large portion of the overall capital cost.

CONNECTIONS BETWEEN INJECTION AND PRODUCTION WELLS Early commercial development focused on techniques used to establish a gas flow path between vertically drilled injection and production wells. One approach was to fracture the coal seam, either with high pressure air or with explosives. Another approach was the reverse combustion technique. High pressure air was injected into the coal seam, and a small stream of air flowed to the production well via naturally occurring cracks and cleats in the coal seam. The coal was ignited at the production well, and burned back towards the injection well. Once the flame front reached the injection well, a cavity between the injection and production wells was established and gasification could proceed. These techniques were problematic. Several attempts at underground coal gasification were abandoned due to a failure to establish connections between the injection and production wells. When the connections were successfully established, flow paths were often less than ideal. The initial flow path tended to follow natural fracture lines in the coal seam, and this gave the operator little control over the initial cavity geometry. Modern directional drilling techniques, developed for oil and gas production, have largely eliminated this problem. Drilling techniques developed for coal bed methane production can be adapted for underground coal gasification. Lightfoot3 has shown that,

Underground Coal Gasiﬁcation

by placing a magnetic target in the production well and a magnetic detector in the drill tip used to drill the injection well, it is possible to connect the two wells even though the horizontal spacing is as much as 1,000 m. In China, UCG operations typically seek to recover fuel from abandoned conventional mines. In the room-and-pillar mining technique, pillars of un-mined coal are left to support the roof. In this case, the initial cavity in the coal bed consists of a mined-out tunnel.

PROCESS CONTROL AND MODELING Compared to the operation of a typical reactor in a chemical plant, the operator of an underground coal gasification cavity has a very limited set of process data. The available data is generally limited to pressure, temperature, gas flow rate, and gas composition at the injection and production wells. The reactor, which consists of the underground cavity, has unknown and constantly changing dimensions. Pirard et al.4 used a helium injection technique to measure cavity volume and back-mixing in an underground coal gasification. On the same test gasification, Brasseur et al.5 used the isotopic exchange of 13 C and 12C to estimate the effective gasification temperature. A number of mathematical models of underground coal gasification have been developed to improve the design and operation of UCG. Recent models include those of Yang6 and Perkins and Sahajwalla.7e9 Yang6 built a large laboratory experiment that simulates the gasification of a thin coal seam. Temperature, pressure, and gas composition measurements in the experiment closely fit model predictions. The model developed by Perkins and Sahajwalla can be divided into two subproblems. The first concerns the description of a small section of the cavity wall.7,8 The second examines the overall changes between the injection and production wells.9 Figure 5.2 qualitatively describes the phenomena at the cavity wall, after oxygen has been depleted by combustion. The wall, proceeding from the bulk gas into the coal seam (left to right in Figure 5.2), can be divided into four zones. The first is a gas film, which describes gas mass transfer to and from the wall using a boundary layer model. The second is an ash layer that covers the wall. The third is a dry zone, and the fourth is the water-soaked coal, called the wet zone. Heat is transferred from the hot bulk gas towards the cool coal seam. The temperature drops to the boiling point of water at the interface between the dry and wet zones. Further into the wet zone, the temperature drops, approaching the unperturbed coal seam temperature. The dry zone can be divided into three sub-zones. Next to the wet zone is the drying sub-zone, where the temperature is above the bulk boiling point of water, but below pyrolysis temperatures. Water in coal micropores is held in place by capillary force, so it has a lower vapor pressure, higher boiling point, than bulk water. This micropore water

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Underground Coal Gasiﬁcation

Gas Film

Ash Layer

Dry Zone

Wet Zone

Heat flux Bulk gas flow Char Gasification

Pyrolysis

CO + H2 ← C + H2O

CO, CO2, H2

2 CO ← C + CO2

H2O, CH4

CH4 ← C + 2 H2

Tar

Drying

Steam

Water influx

Figure 5.2 Qualitative description of phenomena at the UCG cavity wall.7,8

evaporates in the drying zone. In the pyrolysis zone, temperatures are high enough to thermally decompose coal. “Tar” shown in Figure 5.2 represents all C2þ pyrolysis products. In the char gasification sub-zone, char is converted to gas by the steam gasification, CO2 gasification, and direct hydrogenation reactions: o ¼ þ131:46 kJ=gmole CðsÞ þ H2 OðgÞ/COðgÞ þ H2 ðgÞ DHrxn

R-5.1

o ¼ þ172:67 kJ=gmole CðsÞ þ CO2 ðgÞ/2COðgÞ DHrxn

R-5.2

o ¼ 74:94 kJ=gmole CðsÞ þ 2H2 ðgÞ/CH4 ðgÞ DHrxn

R-5.3

Like an above-ground gasifier, the overall rate of gasification is strongly affected by the rate of char gasification. Perkins and Sahajwalla used the kinetic model developed by Roberts and Harris, which was described in Chapter 3. To simplify the calculation, Perkins and Sahajwalla used a power-law approximation of the reaction kinetics, rather than the Langmuir-Hinshelwood form of the rate equations. Unlike a conventional chemical reactor, the UCG cavity has permeable walls. The coal seam is usually saturated with water, and this water is at hydrostatic pressure. Evaporation and pyrolysis create gas. In the char gasification sub-zone, the steam and CO2 gasification reactions (R-5.1 and R-5.2) are faster than the direct hydrogenation reaction (R-5.3), so more gas is produced in this sub-zone than is consumed. Consequently, there is a net bulk flow of gas away from the wall and toward the chamber cavity. The porous solids in the dry zone and the ash layer are a small but significant flow restriction. In a properly operated UCG burn, the cavity pressure will be a little below the hydrostatic pressure.

Underground Coal Gasiﬁcation

The gasification pressure is not an independent design variable. Instead, it is a function of the hydrostatic pressure, which increases with depth of the coal seam. The gasification rate increases with increasing coal depth due to an increase in gasification pressure. Perkins and Sahajwalla8 defined a thermal effectiveness factor, x, as: x ¼ ðchemical energy in product gasÞ=ðheat energy to coal surfaceÞ

Eqn. 5.1

This effectiveness factor increases with increasing depth and pressure for two reasons. The first is that considerable heat is required to evaporate water seeping into the cavity. The latent heat of water vaporization decreases with increasing pressure. The second is that increasing pressure favors the methanation reaction. The reverse of the methane steam reforming reaction is exhibited in Figure 3.3. o COðgÞ þ 3H2 ðgÞ4CH4 ðgÞ þ H2 OðgÞ DHrxn ¼ 206:2 kJ=gmole

R-5.4

The methane formed by this reaction has a high heat of combustion; and the exothermic release of heat reduces the quantity of additional heat required to gasify the coal. The second part of the model by Perkins and Sahajwalla9 is to track the change in gas composition and temperature from the injection well to the production well. This is shown in Figure 5.3. Flammable gases combust in the first ten meters, consuming all of the available oxygen and causing a sharp rise in gas temperature. At the end of the combustion zone, the gases primarily consist of hot, non-flammable N2, CO2, and H2O. In the following 40 meters, steam and CO2 react with char via reactions R-5.1 and R-5.2, so H2 and CO concentrations rise while steam and CO2 levels fall. The methanation reaction, R-5.4, raises the concentration of methane. The gas calorific value increases as the gas flows through the cavity. The temperature falls due to the endothermic nature of the steam and CO2 gasification reactions, R-5.1 and R-5.2, as well as the heat required to evaporate water seeping into the cavity. By the time the gas reaches the production well at 50 meters, the temperature is barely sufficient to sustain a significant rate of gasification. Figure 5.3 shows two reaction zones, a combustion zone in the first 10 meters, and a reduction zone in the following 40 meters. If the cavity were longer, a third zone would appear. In this third zone, the temperature would not be high enough to sustain a significant rate of gasification or pyrolysis; but the gas temperature would continue to fall due to evaporation of water seeping into the cavity. Heat loss due to conduction to the surrounding coal and rock is a relatively small effect for thick coal seams, but this effect can significantly reduce the gas heating value for seams thinner than about 3 meters.10 To simplify the calculations, Perkins and Sahajwalla9 assumed a chamber with a constant, rectangular cross section. Figure 5.3 shows that the temperature falls with increasing distance from the injection well, so one would expect that the cavity would

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Figure 5.3 Gas temperature and composition as a function of distance from the air injection well, from Perkins and Sahajwalla.9

grow quickly near the injection well, and more slowly as one approaches the production well. Kuyper et al.11 predicted a teardrop-shaped cavity for fixed injection and production well sites, with the widest cross section near the injection well and the thinnest cross section near the production well. A CRIP injection scheme, shown in Figure 5.1, with frequent retraction of the injection point, could yield a cavity with a nearly constant cross section. Spalling,7,8 in which chunks of char fall off the wall, can significantly affect the gasification rate. Spalling is difficult to predict. In practice, the UCG cavity will be partially filled with a rubble pile that includes ash, char, and overhead rocks. The gas permeability of this rubble pile is higher than the wall permeability. If the rubble pile completely fills the cavity, then the pile will behave like a packed bed reactor. If the cavity is only partially filled, then much of the bulk gas flow will bypass the rubble pile.

Underground Coal Gasiﬁcation

WATER CONTAMINATION Tests at Hoe Creek in northeastern Wyoming2,12 showed that groundwater contamination is a potential problem. The nature of the problem is illustrated in Figure 5.2. In an above-ground gasifier, coal tars can be avoided by operating at sufficiently high temperatures. In underground coal gasification, regardless of the cavity temperature, a temperature gradient will form at the wall and coal will pyrolyze to form coal tar, a complex carcinogenic mixture. Part of the problem at Hoe Creek was that the cavity pressure was probably too high, which forced some of the gas and tar into the surrounding formation. The test burn was in a shallow seam, and when the roof collapsed, water from a shallow freshwater aquifer mixed with the tar-contaminated coal and rock. Water contamination issues can be reduced by gasifying at slightly less than the hydrostatic pressure. Water will tend to flow into the gasification cavity, and flush coal tars into the gasification zone and towards the production well. This strategy has been successfully demonstrated at the Chinchilla test burn in Australia.13 A low seep rate will provide steam to help gasify the coal. If the pressure is too low, the water flow rate will be excessive; and the heat required to evaporate this excess water will reduce the thermal efficiency of gasification. The coal seam roof can collapse as coal is removed.14,15 This is a common problem in underground mining, but some control techniques, such as roof bolting, are not available when there are no humans underground. If uncontrolled, roof collapse could lead to a loss of gas seal, and provide a path for mixing of contaminated and fresh water. Roof collapse can be controlled, in part, by choosing coal seams with sufficient overburden strength and thickness. In the conventional room and pillar mining technique, pillars of un-mined coal are left to help support the roof. In UCG, the equivalent technique is to limit the cavity cross-sectional area to limit the unsupported roof span, and to space the cavities so that walls remaining between the cavities support the roof. If the UCG seam is much deeper than an overlying fresh water aquifer, there is less likelihood that a roof collapse will lead to significant intermixing of the fresh and contaminated water. Water contamination issues can also be reduced by selecting coal seams that are not hydrogeologically connected with surface waters or water wells. If the water in the coal seam is saline, then this is good evidence that it does not readily mix with fresh water. There will be some local water contamination even if all risk reduction strategies are successfully carried out. When a cavity burn ends, ground water will fill the cavity and mix with the remaining tar. This water will also contain soluble inorganic impurities due to the ash in the char cavity. Fortunately, this contamination does not spread far. Unburned coal tends to absorb compounds from contaminated water; and inorganic rocks, especially clays, have ion-exchange capacity that will buffer the effects of inorganic contamination.

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Underground Coal Gasiﬁcation

UCG-RECOVERABLE COAL Careful seam selection to avoid water contamination and the need to leave unburned coal walls to help support the roof greatly reduce the quantity of recoverable coal. Still, there is a massive quantity of coal that can be recovered by UCG. GasTech16 surveyed the Powder River Basin of Wyoming and Montana to identify coal resources recoverable by UCG. They excluded all coal shallower than 152 m (500 feet) because this coal is feasible to mine using conventional techniques. They also excluded coal deeper than 610 m (2,000 feet) due to the increased cost of recovering deeper coal. Coal seams thinner than 10 m (30 feet) were also rejected. Thinner coal seams can be gasified, but there is little economic incentive to gasify thin seams when thick seams are abundant. A 65% recovery factor was assumed, meaning that 35% of the target seams were left in place for roof support. This reduced the total quantity of UCG-recoverable coal to 200 billion tons. For comparison, North American coal production in 2006 was 1.2 billion tons,17 so, at this rate of consumption, the UCG-recoverable coal from this single basin would last 167 years.

GASTECH PROCESS AND ECONOMIC STUDY GasTech16 completed a process and economic study for UCG in the Powder River Basin.16 They considered two types of UCG operations, one using an air-blown gasification, and the other using an oxygen-blown gasification. Estimates of gas compositions are shown in Table 5.1. The air-blown tests are based on the Hanna tests in south-central Wyoming in the 1970s. The oxygen-blown compositions appear to be Table 5.1 Estimated UCG sygnas compositions used in GasTech study.16 The syngas compostion for the BGL gasiﬁer, fed Powder River Basin coal, is from Bartone and White.18 Gas composition (% dry basis) O2-blown with CO2 removed Component Air-blown O2-blown BGL gasiﬁer

Ar CH4 C2H6 þ C2H4 C3H8 þ C3H10 C4 þ CO CO2 COS H2 H2S N2

0.5 5.4 0.4 0.2 Not reported 16.1 11.8 Not reported 16.7 Not reported 48.8

1.0 10.6 0.8 0.4 Not reported 31.5 23.1 Not reported 32.7 Not reported 0

1.3 13.7 1.0 0.6 Not reported 41.0 0 Not reported 42.5 Not reported 0

0 6.1 0.4 0.1 0.2 56.0 4.4 0.1 30.8 1.2 0.6

Underground Coal Gasiﬁcation

calculated by removing N2 from the air-blown composition. A better estimate could be obtained by applying a model like the one developed by Perkins and Sahajwalla7e9, and using this to determine the effect of changes in feed gas composition. An oxygen-blown gasification would have much higher temperatures near the injection well, but the reduction in gas temperature as the gas moves toward the production well may dampen changes in gas composition (other than the diluting effects of nitrogen) for the air-blown and oxygen-blown cases. In Table 5.1, the air-blown CO level is a little lower, the CO2 level is a little higher, and the H2 level is much higher than the predicted values shown in Figure 5.3. Some of this may be due to pushing the water-gas shift reaction to the right, either by a higher water seepage rate or by lowering the final gasification temperature. For comparison, an estimated syngas composition for a BGL gasifier is also shown in Table 5.1. This gasifier was chosen for comparison to UCG because both the BGL gasifier and UCG operate at low temperatures compared to other types of above-ground gasifiers. The BGL syngas has a relatively high CO concentration, and relatively low levels of CO2 and H2, which are probably due to a low steam feed rate. The UCG syngas shown in Table 5.1 and the BGL syngas all show higher concentrations of hydrocarbons than one would expect from an above-ground gasifier that operates at higher temperatures. These hydrocarbons, especially methane, contribute to the gas fuel value, which is 5.2 kJ/m3 (175 BTU/scf) for the air-blown syngas and 10.2 kJ/m3 (342 BTU/scf) for the oxygen-blown syngas. Coal bed methane can elevate the methane level in UCG syngas, especially for an isolated test cavity. Since the methane concentration for the air-blown syngas in Table 5.1 is very close to the predicted value shown in Figure 5.3, there does not appear to be a significant coal bed methane effect. Syngas hydrocarbons could limit the readily attainable carbon capture level. For example, in a typical IGCC plant with carbon capture and sequestration (see Chapter 9), the syngas is sent to a water-gas shift plant to convert CO and steam to CO2 and hydrogen. The CO2 is then separated from the syngas, compressed, and sequestered, while the hydrogen is burned in a combined cycle plant to produce electric power. Syngas hydrocarbons follow the hydrogen in this process design, so these hydrocarbons would be burned in the gas turbine, producing CO2 emissions. The level of emissions would be much less than a natural gas fired power plant without carbon capture and sequestration (CCS), but it may not be possible to achieve the 90% capture level envisioned for Future Gen. GasTech estimates that the air-blown syngas would cost $1.54/GJ ($1.62/MMBTU) to produce. This gas would be suitable for electric power production without CCS. For power production with CCS, or for the production of liquid fuels or chemicals, the O2-blown syngas is preferable. This gas costs an estimated $2.42/GJ ($2.55/MMBTU). At the time this book was written, this was about half the cost of natural gas.

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Compared to above-ground gasifiers, UCG offers a considerable capital cost savings by eliminating the gasifier, and a considerable operating cost savings by eliminating the cost of mining coal. A disadvantage is that the productivity of a UCG cavity is not constant. If one assumes that the gasification rate per unit of wall area is kept constant, then the gas throughput rate should increase as the cavity grows. This presents challenges to the design and operation of surface facilities, both to prepare the oxygen feed and to process the syngas. Variations in gas throughput can be reduced by using a CRIP design, shown in Figure 5.1, and by simultaneously operating several cavities with staggered ages. UCG syngas shown in Table 5.1 has about a 1:1 ratio of CO to H2, versus about a 2:1 ratio in Figure 5.3. Ratios in this range are suitable for Fischer-Tropsch fluids synthesis with an iron catalyst (see Chapter 13), or dimethyl ether synthesis using the one-pot process (see Chapter 12). Other applications, such as IGCC with CCS, ammonia synthesis, methanol synthesis, and Fischer-Tropsch synthesis with a cobalt catalyst would require a water gas shift reactor to produce higher H2 concentrations (see Chapter 9).

REFERENCES 1. Shafirovich E, Varma A. Underground coal gasification: A brief review of current status. Ind Eng Chem Res 2009;48:7865-7875. 2. Burton E, Friedmann J, Upadhye A. Best practices in underground coal gasification (draft). Lawrence Livermore Nat. Lab., U.S. Dept. of Energy Contract W-7405-7448; 2008. 3. Lightfoot J. Coalbed methane directional drilling and wellhead construction, Underground Coal Gasification Tutorial, 25th Ann. Int. Pittsburgh Coal Conf.; Sept. 29eOct. 2, 2008. 4. Pirard JP, Brasseur A, Coe¨me A, Mostade M, Pirlot P. Results of the tracer tests during the El Tremedal underground coal gasification at great depth. Fuel 2000;79:471-478. 5. Brasseur A, Antenucci D, Bouquegneau J-M, et al. Carbon stable isotope analysis as a tool for tracing temperature during the El Tremedal underground coal gasfication at great depth. Fuel 2002;81:109-117. 6. Yang L. Study on the model experiment and numerical simulation for underground coal gasification. Fuel 2004;83:573-584. 7. Perkins G, Sahajwalla V. A mathematical model for the chemical reaction of a semi-infinite block of coal in underground coal gasification. Energy & Fuels 2005;19:1679-1692. 8. Perkins G, Sahajwalla V. A Numerical study of the effects of operating conditions and coal properties on cavity growth in underground coal gasification. Energy & Fuels 2006;20:596-608. 9. Perkins G, Sahajwalla V. Steady-state model for estimating gas production from underground coal gasification. Energy & Fuels 2008;22:3902-3914. 10. Gregg DW, Edgar TF. Underground coal gasification. AIChE J 1978;24:753-781. 11. Kuyper RA, van der Meer TH, Bruining J. Simulation of underground gasification of thin coal seams. In Situ 1996;20:311-346. 12. Wang FT, Mead SW, Stuermer DH. Water quality monitoring at the Hoe Creek test site: Review and preliminary conclusions. In: Krantz WB, Gunn RD, eds. Underground Coal Gasification: The State of the Art. American Inst Chem Engr Symposium Series; 1983. 13. Linc Energy, UCG and groundwater, <www.lincenergy.com.au/pdf/fact-sheet/e2-fact.pdf; 2010>. 14. Krantz WB, Gunn RD. Modeling the underground coal gasification process: Part IIIdSubsidence. In: Krantz WB, Gunn RD, eds. Underground Coal Gasification: The State of the Art. American Inst Chem Engr Symposium Series; 1983.

Underground Coal Gasiﬁcation

15. Friedmann SJ, Burton E, Buscheck TA, Ezzedine S. Environmental and regulatory aspects, hydrogeology, and simulation, Underground Coal Gasification Tutorial, 25th Ann Int. Pittsburgh Coal Conf; Sept. 29eOct. 2, 2008. 16. GasTech, Inc., Viability of underground coal gasification in the ‘deep coals’ of the Powder River Basin, Wyoming, prepared for the Wyoming Business Council, <www.wyomingbusiness.org/pdf/energy/ WBC_Report_061507_SPM.pdf; 2007>. 17. U.S. Energy Information Administration. <www.eia.doe.gov/fuelcoal.html>. 18. Bartone Jr LM, White J. Industrial Size gasification for syngas, substitute natural gas, and power production. DOE/NETL-401/040607. .

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CHAPTER

6

Sulfur Recovery Contents Coal Combustion Sulfur Compounds in Syngas COS Hydrolysis Water Quench/Water Condensation Acid Gas Removal Processes Physical Solvent: Rectisol Process Physical Solvent: Selexol Chemical Solvents: Amines Chemical Solvents: Benﬁeld Process Chemical Solvents: Aqueous Ammonia Solid Adsorbents for Sulfur Removal Zinc Oxide Regenerable Adsorbents High Temperature Adsorption Elemental Sulfur: Claus Process Shell Claus Offgas Treatment (SCOT) Process Sulfuric and Phosphoric Acid Co-sequestration of CO2 and H2S References

113 114 114 115 116 117 122 124 126 127 128 128 128 129 130 131 132 134 135

Coal, compared to other fossil fuels, has relatively high levels of sulfur. Illinois No. 6 bituminous coal, for example, contains about 2.8 wt.% sulfur on a dry basis. Sulfur exists in coal as organic sulfur compounds and as inorganic pyrites. Crude oil and raw natural gas can also have high sulfur levels, but nearly all of this sulfur is removed before the fuel is used.

COAL COMBUSTION When coal is burned, most of the sulfur is converted to sulfur dioxide (SO2). In the atmosphere, SO2 is gradually oxidized to SO3; which combines with water to form H2SO4, sulfuric acid, the prime component of acid rain. In the USA, SO2 emissions were regulated in the 1960s. This prompted many coalfired power plants to switch to PRB coal, which has only about 0.8 wt.% sulfur. The quantity of PRB coal mined prior to the implementation of SO2 regulations was fairly small. By 2008, the annual production of PRB coal reached 496 million tons, 42% of the US total.1 Coal Gasification and Its Applications. ISBN B978-0-8155-2049-8.10006-3, doi:10.1016/B978-0-8155-2049-8.10006-3

Ó 2011 Elsevier Inc. All rights reserved.

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For newer power plants, switching to a low sulfur coal is not sufficient to meet SO2 regulations. The US Clean Air Act of 19902 sets a sulfur standard at 1.2 lbs SO2/million BTU (0.52 kg SO2/megajoule). To meet this limit, coal combustion plants use limestone scrubbing to capture SO2 from flue gas: SO2 ðgÞ þ CaCO3 ðsÞ/CaSO3 ðsÞ þ CO2 ðgÞ

R-6.1

2CaSO3 ðsÞ þ O2 ðgÞ þ 4H2 Oðg or lÞ/2CaSO4 $2H2 OðsÞ

R-6.2

Sulfur dioxide in the flue gas initially combines with calcium carbonate in the limestone to form calcium sulfite (R-6.1). The calcium sulfite oxidizes to form calcium sulfate; which combines with water to form calcium sulfate dihydrate, better known as gypsum (R-6.2). Coal burning power plants are often accompanied by large stacks of ash and gypsum. If the power plant is next to a coal mine, some of this solid waste may be back-hauled to mined-out areas.

SULFUR COMPOUNDS IN SYNGAS When coal is gasified, the reaction environment is reducing, as opposed to the oxidizing environment in coal combustion. Consequently, sulfur in coal is converted primarily to hydrogen sulfide (H2S) gas, with lesser quantities of carbonyl sulfide (COS). Usually, these sulfur compounds need to be removed from the syngas. Solid catalysts are used to adjust the syngas composition or convert the syngas to chemicals and hydrocarbon fuels, and sulfur compounds are poisons for many of these catalysts. If the syngas is to be burned as a fuel, removal of the sulfur compounds from the syngas will avoid SO2 emissions.

COS HYDROLYSIS Hydrogen sulfide, sulfur dioxide, and carbon dioxide are all considered acid gases, meaning that they form weak acids when dissolved in water. The conventional route to removing H2S and CO2 from syngas is to use a solvent-based acid gas removal process. Carbonyl sulfide, is not an acid gas, so hydrolysis of COS to form H2S is often used to facilitate the removal of the sulfur contained in COS: COS þ H2 O/H2 S þ CO2

R-6.3

An equilibrium constant for this reaction may be written as: K ¼ Where K ¼ equilibrium constant Pi ¼ partial pressure of component i

PH2 S PCO PCOS PH2 O

Eqn. 6.1

Sulfur Recovery

Equilibrium Constant

1,000,000 100,000 10,000 1,000 100 10 0

100

200

300

400

500

600

700

800

900

Temperature, C

Figure 6.1 Effect of temperature on the equilibrium coefﬁcient (Eqn. 6.1) for the COS hydrolysis reaction (R-6.3).

The effect of temperature on the equilibrium constant3 is shown in Figure 6.1. Low temperatures favor the conversion of COS, but the reaction temperature is limited by the kinetics of the reaction and the need to cool the syngas. Several solid catalysts have been used for COS hydrolysis. Undengaard and Berzins4 describe the use of a high surface area alumina (Al2O3), with typical operating temperatures in the 200 to 245 C range, and typical reactor space velocities in the 1,000 to 10,000 h1 range. Kohl and Riesenfeld3 report the use of chromia (Cr2O3)-alumina and copper-chromia-alumina catalysts. The chromia-alumina catalysts typically operate at 315 to 425 C and 250 to 1,000 h1 space velocities. The copper-chromia-alumina catalysts are more active, and typically operate at 260 to 275 C and 2,000 h1 space velocity. Carnell5 reports on the use of a Ni/Mo catalyst on an alumina support. This catalyst is essentially the same as catalysts widely used to hydrogenate petroleum fractions. The hydrolysis of COS is fast, and high reactor space velocities are used. Consequently, COS hydrolysis reactors are small; and the cost of the COS hydrolysis reactor is a small fraction of the overall cost of a gasification-based plant. The primary process consideration for COS hydrolysis is that hot syngas leaving the gasifier must be cooled to achieve favorable equilibrium conditions. Sour gas shift catalysts, sulfur-tolerant water gas shift catalysts, are also active COS hydrolysis catalysts, so a separate COS hydrolysis reactor is not needed. These will be discussed in greater detail in Chapter 7. The sour gas shift reactor conditions are generally set to achieve the desired water gas shift conditions, and COS hydrolysis is then a useful side reaction.

WATER QUENCH/WATER CONDENSATION Direct contact with liquid water is often used to cool and purify syngas. For some gasifiers, such as the Lurgi, BGL, GE, and Siemens gasifiers, a water quench is an integral part of the gasifier design (see Chapter 4). The water from the quench is known as black

115

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syngas cooler syngas

COS hydrolysis

syngas

syngas 40°C

300°C

mercury removal, acid gas removal

sour water

Figure 6.2 Production of sour water when syngas is cooled.

water, because the water contains a wide variety of contaminants that turn the water black. Black water can be recycled to the quench, or, if a slurry feed is used for the gasifier, used as slurry water. If black water must be used for other applications, or discharged, then an elaborate water purification process is required. Relatively clean sour water is produced if solids-free syngas is cooled. This water is called sour (acid) water because it contains dissolved H2S, an acid gas. For conventional mercury removal and acid gas removal processes, syngas must be cooled to near-ambient temperatures. Figure 6.2 shows how sour water may be produced. The sour water typically removes most of the NH3 from the syngas, and a small fraction of the H2S. If the syngas entering the cooler is above about 500 C, then the sour water may also contain significant quantities of halide compounds. Sour water is typically treated using a sour water stripper, a distillation column that produces a mixed NH3/H2S overhead stream and a relatively pure water bottoms stream. If a Claus process is used to produce elemental sulfur, then the mixed gas may be sent to the Claus furnace where NH3 burns to produce N2 and H2O. Alternatively, the Chevron WWT process,6 shown in Figure 6.3, is a two column sour gas stripper configuration used to produce separate NH3 and H2S products. Water has a stronger affinity for NH3 than H2S, and the WWT process uses this difference to separate NH3 and H2S. Ammonia (NH3) can be purified and sold as a byproduct.

ACID GAS REMOVAL PROCESSES A solvent-based acid gas removal (AGR) process is traditionally used to remove nearly all of the H2S and CO2 from syngas. Chapter 10 describes newer, alternative means of separating CO2. All of the AGR processes work in a similar fashion. Syngas is fed to the bottom of an absorption tower and lean solvent, without dissolved acid gas; is fed to the top of the tower. Cleaned syngas leaves the top of the tower, and rich solvent, solvent with dissolved acid gas, leaves the bottom of the tower. The rich solvent is then depressurized, and, or heated to free the acid gas and regenerate the lean solvent. AGR processes may be divided into two categories, those that use a physical solvent, and chemical solvents that react with the acid gas. Physical solvents tend to have a weaker

Sulfur Recovery

H2S stripped water H2S stripper

NH3

steam recycle water degasser NH3 stripper sour water

recycle water

steam

feed/effluent heat exchanger

stripped water

Figure 6.3 The Chevron WWT process,6 a two column sour water stripper process used to produce separate H2S and NH3 product streams.

affinity for acid gasses, which means that the lean solvent does not rapidly absorb the acid gas; but the rich solvent may be regenerated with a relatively small energy input. Chemical solvents, on the other hand, more strongly bind to the acid gases, which are useful when the acid gas components have low partial pressures; but the rich chemical solvents require more energy to regenerate.

PHYSICAL SOLVENT: RECTISOL PROCESS The Rectisol process is one of the oldest commercial AGR processes still in use. This process was developed by Lurgi to process syngas from the Lurgi gasifier. In addition to the usual syngas components, Lurgi syngas contains low molecular weight hydrocarbons due to the low gasifier operating temperature. Refer to Table 4.1, which gives syngas compositions for the similar BGL gasifier. In the Rectisol process, the solvent is chilled methanol. Table 6.1 shows the solubility of syngas compounds, relative to CO2, in methanol at 25 C. Rectisol separates these compounds based on their relative solubility. As a first approximation, known as Henry’s Law, the solubility of gas compounds increases linearly with respect to partial pressure. There are numerous variations of the Rectisol process, each modified to achieve specific process objectives. The version shown here, in Figure 6.4, is from Ranke and

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Sulfur Recovery

Table 6.1 Relative solubility7 of syngas compounds in methanol at 25 C. Compound Relative solubility

Benzene CH4 C2H4 C2H6 C3H8 CO CO2 COS H2 H2S N2

59.5 0.05 0.46 0.42 2.35 0.02 1) 3.92 0.005 7.06 0.01

)

Basis of relative solubility.

Mohr7, and is designed to process syngas from a sour gas shift reactor prior to ammonia synthesis. The ammonia plant produces 1,000 tons/day. A portion of the ammonia is to be converted into urea, (NH2CONH2), so a relatively pure CO2 stream is recovered as a feedstock for the urea plant. As shown in Table 6.2, the feed gas has relatively high levels of H2 and CO2, and a relatively low level of CO, which is the result of the sour gas shift. The water content of the feed gas is not shown, but the feed gas contains significant quantities of water; and the Rectisol plant is designed to remove water from the syngas. A portion of the Rectisol plant is shown in Figure 6.4a. Syngas is mixed with methanol and chilled by exchanging the feed syngas/methanol mixture with cold product gases. The gas/liquid mixture is then separated in a flash drum. The liquid phase is a methanol/water mixture and contains most of the water from the syngas feed. This liquid is depressurized, warmed, and then fed to the water splitter, a binary distillation column. Steam is used to heat the reboiler. Water, as the higher boiling component, is withdrawn from the bottom of the column. Instead of a condenser, liquid methanol is fed to the top of the column, and methanol vapor is withdrawn as the overhead product. Gas from the flash drum is fed to the primary column, a single absorption tower split into three sections. This column is designed to selectively remove H2S and CO2 based on their relative solubilities. The bottom section removes H2S and COS using CO2-loaded methanol from the middle section. Gas rising out of the bottom section and into the middle section is nearly H2S-free. The top and middle sections remove CO2 from the rising syngas. Chilled methanol is fed to the top section. Liquid from the top section is withdrawn, cooled by exchange with another process stream, and reinjected into the middle section. Refrigeration coils in the middle section remove the heat of absorption. Purified syngas, with low levels of H2S and CO2, leaves the top of the column.

Sulfur Recovery

Purified syngas

chilled methanol primary column

refrigerant

methanol + CO2

cold product gas syngas feed

methanol + H2S

methanol warm product gas

flash gas compressor methanol vapor liquid methanol

water splitter

steam

water

Figure 6.4a Portion of a Rectisol plant7 designed to process syngas prior to ammonia synthesis and subsequent conversion of ammonia and CO2 to urea.

Liquids are withdrawn from the middle and bottom sections. Both liquid streams are depressurized; and the gasses flashing off the liquids are recompressed and reinjected into the bottom of the primary column. Liquid withdrawn from the middle section has high levels of dissolved CO2, and very little dissolved H2S. Liquid withdrawn from the bottom of the column contains nearly all of the H2S and COS contained in the feed syngas. The remainder of the Rectisol plant is shown in Figure 6.4b. The CO2 and tail gas columns are designed to remove most of the dissolved CO2 without also removing

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Sulfur Recovery

H2S tail gas

CO2

cooling water

methanol + CO2

methanol vapor

methanol + H2S H2S stripper

steam tail gas column

CO2 column N2

methanol

Figure 6.4b Remainder of the Rectisol plant.

Table 6.2 Feed and product gas compositions,7 utilities for the Rectisol plant shown in Figure 6.4. Component, Feed Puriﬁed Tail gas dry basis syngas syngas CO2

H2 N2 þ Ar CO þ CH4 CO2 H2S þ COS Flowrate, kmole/h Pressure, MPa Power, shaft Steam, 0.5 MPa Refrigeration Cooling water Nitrogen at 0.3 MPa

H2S

62.47% 0.51% 2.67% 34.10% 0.25% 6,021.1

95.27% 0.76% 3.97% 20 ppm 0.1 ppm 3,936.1

0.84% 0.08% 0.33% 98.75% 2 ppm 1,283.9

0.07% 25.32% 0.04% 74.57% 5 ppm 1,003.6

0.31% 4.36% 0.08% 68.01% 27.24% 54.2

7.8

7.6

0.18

0.11

0.25

1,100 6,490 1,980 5,000 256.7

kW kg/h kW lpm kmole/h

Sulfur Recovery

dissolved H2S and COS. The CO2 column is a three-section absorption tower. Methanol plus CO2 from the primary column is fed to the top of the top section. Methanol plus H2S is fed to the top of the middle section. Liquid is withdrawn from the bottom of the middle section, warmed, and reinjected into the bottom of the bottom section. Warming this liquid causes part of the dissolved gas to boil off, which establishes a vapor flow in the bottom of the CO2 column. As the gas flows upward in this column, downward flowing liquid absorbs H2S and COS in this gas, so the CO2 leaving the top of the column is nearly sulfur free. Vapor flow in the bottom of the tail gas column is established by injecting nitrogen into the column. Since the gasifier is equipped with an air separation unit to supply O2, an abundance of nitrogen is available. Tail gas, largely a blend of CO2 and N2, is withdrawn from the top of the column. Liquid from the top section of the CO2 column is withdrawn and split. A portion is reinjected into the top of the middle section of the CO2 column, and the remainder is fed to the top of the tail gas column. Liquid from the bottom of the top section of the tail gas column is withdrawn and injected into the top of the bottom section of the CO2 column. Liquid from the bottom of the tail gas column is warmed and fed to the H2S stripper. Methanol vapor, from the water splitter shown in Figure 6.4a, is also fed to the H2S stripper. Vapor from the top of the stripper is cooled, producing the H2S gas product and condensing methanol, which is returned to column. As shown in Table 6.2, the H2S gas product contains more CO2 than H2S, but the H2S content of this stream is high enough to feed to a Claus furnace for the production of elemental sulfur. The reboiler is steamheated. The bottom product is methanol, free of dissolved gasses. The tail gas, which is vented to the atmosphere, contains 36% of the CO2 in the feed syngas. This process was designed before carbon capture and sequestration was considered a priority. The only reason that a relatively pure CO2 stream is produced is that it is a feedstock for urea production. For a gasification system with carbon capture and sequestration, the system would be redesigned to prevent venting this tail gas stream. For this particular application, ammonia synthesis, the tail gas could be recycled to the primary column. This would put N2 in the purified gas, but since N2 is a reactant in ammonia synthesis, this is acceptable. Ranke and Mohr7 also describe a simpler version of the Rectisol process that does not produce separate CO2 and H2S streams. Lurgi8 shows a Rectisol process designed to process syngas for methanol synthesis. In this version of the process, syngas is sent to the first half of the Rectisol process to remove sulfur compounds. The sulfur-free syngas is then sent to a water gas shift reactor to adjust the H2/CO level to the ratio desired for methanol synthesis. The shifted gas is then sent to the back half of the Rectisol process, where CO2 is removed. The original Rectisol plant, part of the Great Plains Synfuels plant, was designed to remove light hydrocarbons from the

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Sulfur Recovery

syngas produced by their Lurgi gasifiers. Naphtha, gasoline boiling range hydrocarbons are produced as a byproduct. An advantage of the Rectisol process is that the low operating temperatures tend to trap trace impurities in the syngas, which reduces catalyst poisoning in downstream reactors. The ultimate fate of these trapped impurities is not clear. Methanol, compared to other AGR solvents, is inexpensive. On the other hand, the complexity of the Rectisol process and the extensive use of refrigeration makes this an expensive AGR process, in terms of both capital and operating costs. Newer gasification plant designs generally use less expensive AGR processes.

PHYSICAL SOLVENT: SELEXOL There are a number of AGR processes that use a polar organic solvent to physically dissolve H2S and CO2. One of the more popular processes is Selexol, which uses a madefor-purpose solvent consisting of dimethyl ethers of polyethylene glycol oligimers: CH3 O-ðCH2 CH2 OÞn -CH3 The oligimer number, n, is between 3 and 9. The selective removal of H2S and CO2, like Rectisol, is based on the higher relative solubility of H2S in the Selexol solvent compared to CO2. Table 6.3 shows relative gas solubilities in Selexol solvent. A flowsheet for the selective removal of H2S and CO2 using the Selexol process is shown in Figure 6.5. Syngas is fed to the bottom of the sulfur absorber. The solvent fed to the top of this column is CO2-loaded Selexol. Because the solvent has a greater affinity for H2S than for CO2, this column absorbs H2S but very little CO2. The H2S-depleted syngas is fed to the CO2 absorber, where CO2 is removed by lean Selexol. The CO2 affinity of the Selexol solvent is sufficiently weak; so the bulk of the CO2 can be removed by simply depressurizing the solvent. Depressurization is done in two steps. Gas flashed during the first pressure let-down is recompressed and reinjected into the CO2 absorber. Gas flashed during the second depressurization is the CO2 product stream. 9

Table 6.3 Relative solubilities of gases in Selexol solvent,9 Compound Relative Solubility

H2 ) N2 CH4 CO2 COS H2S H2O )

1.0 1.5 5 76 175 670 55,000

Basis for relative solubility.

Sulfur Recovery

purified syngas H2S

CO2

CO2 absorber

sulfur absorber

lean solvent

CO2 loaded solvent

syngas

H2S loaded solvent

water H2S stripper

Figure 6.5 Selexol process10,11 for the selective removal of H2S and CO2.

The Selexol solvent creates a strong bond to H2S, so a distillation column is used to strip H2S from the solvent. The H2S loaded solvent from the bottom of the H2S absorber is warmed by heat exchange with the hot, lean solvent. This is produced by the H2S stripper. The flashed gases are recompressed and reinjected into the H2S absorber. The liquid is fed to the H2S stripper. The overhead gas produced from the H2S stripper is an H2S rich gas product. The overhead condensed liquid is water. This water is added or removed from the condensate return stream as needed to adjust the water content of the solvent. The bottom product from the stripper is the lean Selexol solvent, which is cooled and sent to the top of the CO2 absorber.

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Sulfur Recovery

Selexol plants may also be configured to selectively remove H2S, or non-selectively remove both H2S and CO2. If the CO2 absorber in Figure 6.5 is eliminated, then the resulting process, at low solvent/gas ratios, will preferentially remove H2S because of the relatively high solubility of H2S. Much or most of the CO2 remains in the treated gas. The same flowsheet, with a higher solvent/gas ratio, non-selectively removes both H2S and CO2.

CHEMICAL SOLVENTS: AMINES A wide variety of AGR processes rely on the reactions of organic amine compounds with acid gases to form weak compounds. These compounds are decomposed by heat to release the acid gas and regenerate the solvent. The use of mono-ethanol-amine (MEA) to remove CO2 from flue gas was shown in Figure 2.5. Figure 6.6 shows amine compounds commonly used in AGR processes. In addition to these compounds, other amines and proprietary formulations, that include corrosion inhibitors and defoaming agents, are also used. The amines typically include one or more alcohol functional groups to improve water solubility and decrease vapor pressure. A low molecular weight amine, such as MEA, reduces the mass of solvent required. On the downside, low molecular weight amines also have higher vapor pressures which leads to greater solvent loss. Some MEA AGR processes use a water wash at the top of the absorption or stripping towers to reduce MEA losses. The amine group reacts with the acid gas. Simple acid-base reactions include: RNH2 ðaqÞ þ H2 SðgÞ4RNH2 ðaqÞ þ Hþ ðaqÞ þ HS ðaqÞ4RNHþ 3 ðaqÞ þ HS ðaqÞ R-6.4

RNH2 ðaqÞ þ CO2 ðgÞ þ H2 OðlÞ4RNH2 ðaqÞ þ Hþ ðaqÞ þ þ HCO 3 ðaqÞ4RNH3 ðaqÞ þ HCO3 ðaqÞ

HH

H

HO-C-C-N H H

H H HO- C-C

H

monoethanolamine (MEA) H H HO HHO H H-C-C-C-C-C-C-H HH N HH H diisopropanolamine (DIPA)

H H

HH N H

C-C-OH HH

diethanolamine (DEA) H H HO- C-C H H

HH N

C-C-OH HH

C H

H H Methyldiethanolamine (MDEA)

Figure 6.6 Amines commonly used in acid gas removal processes.

R-6.5

Sulfur Recovery

R in these reactions is an unspecified organic functional group. The reactions are written as if the amine is a primary amine, but secondary and tertiary amines will also react in this manner. The reaction of CO2 with water to form carbonic acid (H2CO3) is slow. A faster reaction, which prevails at lower CO2/amine molar ratios, is the reaction of a mole of CO2 with two moles of an amine to form a carbamate compound: 2RNH2 ðaqÞ þ CO2 ðgÞ4RNHCOONH3 RðaqÞ

R-6.6

The reaction is written as if the amine is a primary amine. Both primary and secondary amines undergo the carbamate reaction, but tertiary amines do not. Maddox et al.12 show that the carbamate compound is the primary reaction product for primary and secondary amines and CO2 up to a CO2/amine molar ratio of about 0.5, the stoichiometric ratio in reaction R-6.6. To get higher CO2/amine loadings, the acid/base reaction, R-6.5, with a 1:1 stoichiometric ratio, is used. Tertiary amines such as MDEA do not undergo the carbamate reaction, so they have a lower affinity for CO2 and a higher H2S/CO2 selectivity. Table 6.4 shows typical operating conditions for amine AGR processes.13 The acid gas/amine molar ratio is typically kept below the 0.5 stoichiometric ratio for the carbamate reaction. This is so that carbon steel wetted surfaces can be used. Higher acid gas/amine ratios can be used, but the solution changes from alkaline to acidic and stainless steel is then required for corrosion control. With a higher acid gas/amine ratio, the quantity of solvent is reduced. This reduces the heat required to strip the solvent, but increases capital costs due to the use of more expensive materials. Hindered amines14 are primary or secondary amines with bulky organic groups near the amine functional group. Carbon dioxide is rapidly absorbed by these compounds to form carbamates, but the carbamate compound is unstable due to stearic hindrance by the bulky organic group. Consequently, the carbamate may decompose, allowing the CO2 to be absorbed in the acid/base form shown in reaction R-6.5. This approach combines the rapid absorption rate of the carbamate mechanism (R-6.6) with the high

Table 6.4 Typical operating conditions for amine-based acid gas removal processes.13

Solution strength, wt.% Acid gas loading, mole/mole amine DHrxn for H2S, KJ/mole) DHrxn for CO2, KJ/mole) )

MEA

Amine DEA

MDEA

15e20 0.3e0.4

25e35 0.3e0.4

30e50 0.7e0.9

41 84

38 67

39 61

Heat of reaction for acid gas loadings less than 0.5 mole acid gas/mole amine.

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Sulfur Recovery

solvent loading of the acid/base mechanism (R-6.5). This is the basis of the ExxonMobil Flexsorb process.15

CHEMICAL SOLVENTS: BENFIELD PROCESS Benson, Field, and coworkers designed the Benfield process16e17 to remove H2S and CO2 from coal-derived syngas prior to ammonia synthesis. This process was described in Chapter 2 as a component of the Sargas process. This combines pressurized fluidized bed combustion of coal with the Benfield process for carbon capture. The Benfield process and the similar Catacarb process18 rely on the reversible reactions of acid gases with a potassium carbonate solution: K2 CO3 ðsÞ þ nH2 O/2Kþ ðaqÞ þ CO32 ðaqÞ þ nH2 O

R-6.7

CO32 ðaqÞ þ CO2 ðgÞ þ H2 OðlÞ42HCO3 ðaqÞ

R-6.8

CO32 ðaqÞ þ H2 SðgÞ4HCO3 ðaqÞ þ HS ðaqÞ

R-6.9

Sodium carbonate, in concept, could be used in place of potassium carbonate. POTASSIUM carbonate is preferred because potassium bicarbonate (KHCO3) has a much higher water solubility than sodium bicarbonate (NaHCO3). In the original process, the intent was to vent CO2 to the atmosphere. Since the partial pressure of CO2 in the syngas was higher than 1 atmosphere (0.1 MPa), the difference in CO2 partial pressure was used to move reaction R-6.8 to the right in the absorption tower and to the left in the stripping tower. Both towers operate at a little above 100 C. Since reactions R-6.8 and R-6.9 are exothermic, steam is used in the stripper reboiler to add heat to the solution to maintain the solution temperature. This will offset the endothermic heat of the reverse reactions. The reaction of CO2 with water to form bicarbonate ions is slow. Refer to the discussion following reaction R-6.5 with regard to bicarbonate formation in amine absorption. Consequently, a wide range of promoters are used to increase the rate of absorption. Amines19e22 are the most commonly used promoters, but inorganic promoters23,24 have also been used. Astarita et al.25 noted that all of the promoters work in a similar fashion. A weak compound rapidly forms between CO2 and the promoter, which will gradually release CO2 to form the bicarbonate ion. This is shown in reaction R-6.8. Danckwerts and Sharma19 used penetration theory to describe the absorption as a coupled kinetic and mass transfer process using the following equation: qﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃ Eqn. 6.2 Nco2 ¼ AðCi Cb Þ ko2 L þ Dco2 kOH ½OH þ Dco2 kb ½prom Where: Nco2 ¼ absorption rate of CO2, gmol/s

Sulfur Recovery

A ¼ gaseliquid interfacial surface area, cm2 Ci ¼ concentration of unreacted CO2 at the gaseliquid interface, gmole/L Cb ¼ concentration of unreacted CO2 in the liquid bulk, gmole/L k0L ¼ physical mass transfer coefficient, cm/min Dco2 ¼ diffusivity of CO2, cm2/s kOH ¼ rate constant for the reaction between CO2 and hydroxyl ions, L/gmole-s [OH] ¼ concentration of hydroxyl ions, gmole/L kb ¼ rate constant for the reaction between CO2 and the promoter, L/gmol-s [Prom] ¼ concentration of promoter, gmole/L. Cents et al.26 re-examined the derivation of this equation to justify the use of an approximation used by Dankwerts and Sharma.

CHEMICAL SOLVENTS: AQUEOUS AMMONIA Aqueous ammonia can be used to remove acid gases from syngas. A challenge in designing this process is that ammonia has a high vapor pressure; so the process must be designed to limit ammonia loss. Kohl and Reisenfeld3 describe the use of aqueous ammonia to remove H2S and CO2 from coal-derived gases, primarily coke oven gas. They show that the vapor pressure of ammonia falls with increasing dissolution of H2S and CO2 due to the formation of ammonium salts. Dissolution rates for H2S are much higher than CO2, so the absorption can be somewhat selective for H2S by using short contact times. Powerspan27,28 markets aqueous ammonia processes for SO2 and SO2 plus CO2 removal from combustion flue gas. In the SO2 only process, ammonia is dissolved in a recycle ammonium sulfate solution: R-6.10 NH3 ðgÞ þ H2 OðlÞ4NHþ 4 ðaq þ OH ðaqÞ Sulfur dioxide is then absorbed using this solution. Hydroxyl ions in the scrubbing solution react with SO2 to form sulfite ions: SO2 ðgÞ þ H2 OðlÞ42Hþ ðaqÞ þ SO2 3 ðaqÞ

R-6.11

Hþ ðaqÞ þ OH1 4H2 OðlÞ

R-6.12

The sulfite ions are then oxidized to sulfate ions by blowing air through the solution. Air not consumed by the reaction mixes with the flue gas and goes up the flue stack. O2 ðgÞ þ 2SO32 42SO42

R-6.13

Combined, reactions R-6.10 to R-6.13 give: 4NH3 ðgÞ þ 2SO2 ðgÞ þ 2O2 ðgÞ44NH4 þ ðaqÞ þ SO42 ðaqÞ

R-6.14

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Sulfur Recovery

The ammonium sulfate product can be sold as a fertilizer, either as a solution or as solid crystals. The SO2 plus CO2 process uses two absorption towers. The first tower operates in a similar fashion to the SO2 only process. The second tower strips CO2 from the SO2free flue gas. The ammonium carbonate solution from the second tower is stripped using steam to regenerate the ammonium hydroxide solution and free CO2 gas. In both processes, flue gas is scrubbed with water just before it is vented to the stack to recover gas phase ammonia. The Powerspan process relies on water evaporation by hot flue gas during SO2 absorption to balance water use across the plant. Adaption of this process to H2S and CO2 removal from syngas would require modification of the sulfite oxidation step, (R-6.13) and a re-examination of the process water balance.

SOLID ADSORBENTS FOR SULFUR REMOVAL For integrated gasification combined cycle (IGCC) plants, the liquid solvent based AGR processes are generally sufficient to meet sulfur emission standards. For other applications, the syngas will be processed in catalytic reactors, and many of these catalysts are poisoned by sulfur. To obtain reasonable catalyst life, sulfur concentrations are further reduced to very low levels using solid adsorbents.

Zinc Oxide The traditional adsorbent used for removing low levels of H2S from syngas is zinc oxide (ZnO), which reacts with H2S to form zinc sulfide and water: ZnOðsÞ þ H2 SðgÞ/ZnSðsÞ þ H2 OðgÞ

R-6.15

Zinc sulfide forms on the surface of the adsorbent, and diffuses into the bulk.5 Increasing the adsorption temperature from about 100 C to about 400 C increases the sulfur adsorption capacity from about 30% to about 90% of the theoretical maximum due to the activation energy for the solid diffusion mechanism. Since ZnS has a higher molar volume than ZnO, there is a loss of porosity in the adsorbent. The bulk of the adsorbent is ZnO, but additional compounds are used to produce a solid with a high surface area and a reasonable physical strength. Traditionally, the adsorbent is not regenerated. Consequently, this adsorbent is not practical for use with high H2S concentrations. The feed gas to a ZnO bed is normally a natural gas-derived syngas, or a coal-derived syngas that has had the bulk of the sulfur removed by a liquid solvent-based AGR process.

Regenerable Adsorbents A regenerable solid adsorbent should, like the traditional ZnO adsorbent, reduce the H2S concentration to very low levels, but would be more practical with higher concentrations

Sulfur Recovery

of H2S in the feed gas. Yi et al.29 describe the regeneration of a proprietary, ZnO-based adsorbent. Oxygen converts the ZnS back to ZnO and produces sulfur as SO2: 2ZnSðsÞ þ 3O2 ðgÞ/2ZnOðsÞ þ 2SO2 ðgÞ

R-6.16

RTI30 markets a similar sulfur capture process that uses two fluidized beds, one for the sulfur capture, reaction R-6.15, and the other for the adsorbent regeneration, reaction R-6.16. Sulfur capture rates of 99.8 to 99.9% are claimed. Kobayashi et al.31 and Shirai et al.32 describe a similar regenerable sulfur adsorbent based on zinc ferrite.

High Temperature Adsorption A typical AGR process requires cooling the syngas to near ambient temperatures. High temperature syngas treatment could eliminate, or at least reduce, the need for syngas cooling; and this could improve the overall thermal efficiency of the process. Cheah et al.33 reviewed mid to high temperature solid sulfur sorbents, and focused on sorbents based on zinc, copper, calcium, manganese, iron, and rare earths. Only a few of the 179 references that they cited will be described here. Lew et al.34 described sulfur adsorption with zinc titanate at 650 C. The titanate component prevents reduction of zinc to metal in syngas. Sulfur is captured via the following reactions: Zn2 TiO4 ðsÞ þ 2H2 SðgÞ/2ZnSðsÞ þ TiO2 ðsÞ þ 2H2 OðgÞ ZnTiO3 ðsÞ þ H2 SðgÞ/ZnSðsÞ þ TiO2 ðsÞ þ H2 OðgÞ

R-6.17 R-6.18

The adsorbent is then regenerated with oxygen, producing sulfur dioxide: 2ZnSðsÞ þ TiO2 ðsÞ þ 3O2 ðgÞ/Zn2 TiO4 ðsÞ þ 2SO2 ðgÞ 2ZnSðsÞ þ 2TiO2 ðsÞ þ 3O2 ðgÞ/2ZnTiO3 ðsÞ þ 2SO2 ðgÞ

R-6.19 R-6.20

Wakker et al. used MnO or FeO supported on g-alumina. Sulfur was adsorbed via the following reversible reaction, where Me represents either manganese or iron: 35

MeOðsÞ þ H2 SðgÞ4MeSðsÞ þ H2 OðgÞ

R-6.21

The adsorbent was regenerated with steam, via the reverse reaction. Wakker et al. tested the adsorption between 400 and 800 C, and concluded that the optimum between adsorbent capacity and stability was at about 600oC. The iron-based sorbent was tested for about 110 cycles, and the manganese-based sorbent was tested for about 310 cycles. Sulfur capacity declined quickly during the first ten cycles, and then more gradually for the remainder of the test cycle.

129

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Sulfur Recovery

Fenouil and Lynn36,37 measured H2S adsorption using both calcined limestone (CaO) and un-calcined limestone (CaCO3). A 560 to 860 C temperature range was used for the un-calcined limestone, which reacts as follows: CaCO3 ðsÞ þ H2 SðgÞ/CaSðsÞ þ CO2 ðgÞ þ H2 OðgÞ

R-6.22

The conversion of CaCO3 to CaS was limited to about 10% because of sintering of the product layer. At about 950 C and atmospheric pressure, limestone is completely calcined: CaCO3 ðsÞ/CaOðsÞ þ CO2 ðgÞ

R-6.23

The calcined limestone may then be used to capture sulfur via the following reaction: CaOðsÞ þ H2 SðgÞ/CaSðsÞ þ H2 OðgÞ

R-6.24

Fenouil and Lynn showed that CaO could be completely converted to CaS. They tested a temperature range of 560 to 1,100 C. The equilibrium H2S concentration increases above CaO with increasing temperature. At the Process Development Facility operated by the Southern Company in Wilsonville, Alabama, there was an attempt to capture sulfur in the Transport gasifier by cofeeding limestone and coal. They found that sulfur capture was ineffective because the gasification temperatures were too high to thermodynamically favor CaS formation. This approach may work if the product syngas were to be cooled before fly ash removal. As mentioned in Chapter 3, calcium is also a gasification catalyst; so there is the possibility that calcium could serve as both a catalyst and as a sulfur sorbent. Calcium sorbents are not expected to regenerated. Limestone is a very abundant and inexpensive material, so there is little financial incentive to recover calcium. If sulfur levels in the syngas are low and sulfur prices are low, then limestone capture of sulfur may be economically attractive.

ELEMENTAL SULFUR: CLAUS PROCESS After removing H2S from syngas, the H2S is usually converted to elemental sulfur via the Claus process,38 shown in Figure 6.7. A gas stream containing H2S is partially burned in air to convert 1/3 of the H2S to SO2: H2 SðgÞ þ 2O2 ðgÞ/SO2 ðgÞ þ H2 OðgÞ

R-6.25

The SO2 reacts with the unburned H2S to form elemental sulfur and water: SO2 ðgÞ þ 2H2 SðgÞ/3SðgÞ þ 2H2 OðgÞ

R-6.26

The air rate to the burner is controlled so that only 1/3 of the H2S burns, thereby maintaining the proper stoichiometry for reaction R-6.26.

Sulfur Recovery

steam H2S

condenser

gas

burner

water sulfur

air

reheat

reheat

converter

converter

reheat

tail gas

condenser sulfur

converter

condenser sulfur

condenser sulfur

Figure 6.7 Three-stage Claus process for the conversion of H2S to elemental sulfur.

Hydrogen sulfide and sulfur dioxide are partially converted in the burner to elemental sulfur and water. A boiler, immediately following the burner, uses some of the heat of combustion to generate steam. A condenser then further cools the gas to condense liquid sulfur. Figure 6.7 shows a typical configuration of a three-stage Claus plant. This plant contains three catalytic converters. Each converter pushes reaction R-6.26 further toward completion. A reheater prior to each converter raises the gas mixture to a few degrees above the sulfur dew point. A condenser after each converter condenses the sulfur. For the burner shown in Figure 6.7 to work properly, the H2S content of the feed gas needs to be above about 55 to 60%. At lower H2S concentrations, a split flow scheme is often used. Much of the feed gas bypasses the burner, steam generator and first condenser. Most of the H2S fed to the burner is converted to SO2, and the cooled burner exhaust is mixed with bypass gas to get the desired 2/1 H2S/SO2 ratio. If part or all of the air is replaced by oxygen, the burner can treat feed streams with still lower H2S concentrations.

SHELL CLAUS OFFGAS TREATMENT (SCOT) PROCESS The tail gas shown in Figure 6.7 consists mostly of nitrogen and argon, the inert components in air fed to the burner. There is also a substantial quantity of H2S and SO2,

131

132

Sulfur Recovery

however, the overall conversion in a Claus plant is typically no more than 97%. These sulfur gases are toxic, controlled air pollutants, so the tail gas cannot be directly discharged to the atmosphere. The SCOT process is used to capture sulfur gases in Claus tail gas. Hydrogen and tail gas are fed to a catalytic hydrogenation reactor to convert all sulfur species to H2S. This gas is then sent to an acid gas removal process based on MDEA. The sulfur-free tail gas is discharged to the atmosphere; and the captured H2S is recycled to the Claus feed stream.

SULFURIC AND PHOSPHORIC ACID In 2000, the USA consumed 13.3 million tons of sulfur or sulfur equivalent.39 Of this quantity, 1.02 million tons (7.7%) was sulfuric acid produced as a byproduct of smelting pyrite ores, and another 0.46 million tons (3.5%) was imported sulfuric acid. The remaining 88.8% of sulfur consumed was elemental sulfur. Most elemental sulfur produced in 2000 was a byproduct of natural gas and crude oil processing. With an increase in coal gasification, sulfur from coal should become a larger component of sulfur supply. Natural deposits of elemental sulfur are produced using the Frasch process. Rising volumes of byproduct sulfur depressed sulfur prices, and Frasch sulfur declined to less than 7% of US production by 2000. Nearly all elemental sulfur is converted to sulfuric acid (H2SO4). Elemental sulfur is burned with air in a furnace to produce sulfur dioxide: SðlÞ þ O2 ðgÞ/SO2 ðgÞ

R-6.27

The hot gases are then passed over solid catalyst, usually based on V2O5, to further oxidize the sulfur dioxide. 3SO2 ðgÞ þ O2 ðgÞ/2SO3 ðgÞ

R-6.28

Sulfur trioxide reacts with water in an absorber to produce sulfuric acid. SO3 ðgÞ þ H2 OðlÞ/H2 SO4 ðlÞ

R-6.29

Sulfuric acid is a basic chemical with a wide range of uses. The phosphate fertilizer uses 34% of the sulfuric acid consumed each year. Chemical fertilizers are generally classified according to their NPK content, representing the concentrations of the three principal plant nutrients, nitrogen (wt.%), phosphorus (wt.% P2O5 equivalent), and potassium (wt.% K2O equivalent). Sulfur is a secondary nutrient, so less is required than the primary nutrients; but more than the micronutrients, such as copper, which are needed but at very low concentrations. Phosphate ore contains 65e75 wt.% fluorapatite (Ca10(PO4)6F2). The balance consists of impurities, including SiO2. Phosphate ore is ground and mixed with recycled

Sulfur Recovery

dilute phosphoric acid and sulfuric acid to produce phosphoric acid, waste solid gypsum (CaSO4$2H2O), and a silicon tetrafluoride byproduct. 2Ca10 ðPO4 Þ6 F2 ðsÞ þ 20H2 SO4 ðaqÞ þ SiO2 ðsÞ þ 18H2 OðlÞ/12H3 PO4 ðaqÞ þ SiF4 ðgÞ þ 20CaSO4 $2H2 OðsÞ R-6.30 The phosphate ore typically contains small quantities of limestone impurities, which react with sulfuric acid to form additional gypsum, soluble magnesium ions, and CO2 gas. Gases from the process are scrubbed with water to recover hexaflourosilicic acid (H2SiF6), which is a valuable byproduct of the gasification process. 3SiF4 ðgÞ þ 2H2 OðlÞ/2H2 SiF6 ðaqÞ þ SiO2 ðsÞ

R-6.31

The phosphoric acid/gypsum slurry is filtered, and the solid gypsum is sent to a waste stack. The phosphoric acid is sent to evaporators for concentration, water removal, and is then converted to phosphate fertilizers, such as ammonium phosphates. It is important to note, that over half of the sulfur produced in the USA ultimately resides in a waste gypsum stack next to a phosphoric acid plant. For comparison, when coal is burned in a pulverized coal power plant, most of the sulfur in the coal also ends up in a waste gypsum stack, as the product of flue gas desulfurization. There is a small commercial market for gypsum in wallboard and cement manufacture, but large quantities of relatively pure natural gypsum are available and relatively inexpensive. Waste gypsum produced by power plants and phosphoric acid plants is too impure to use without costly processing. Most phosphoric acid producers prefer to buy elemental sulfur and make their own sulfuric acid, rather than purchase sulfuric acid. The main reason for this is due to the highly exothermic sulfuric acid process generates excess steam, which can be used in the phosphoric acid evaporators. A secondary reason is that the atomic weight of sulfur is less than the molar weight of H2SO4. A typical phosphoric acid plant produces about 1,000 tons/day of P2O5 equivalent, and requires either 915 tons/day of elemental sulfur or 2,800 tons/day of sulfuric acid for production. The smaller mass of the elemental sulfur supply makes it less expensive to ship. An alternative to producing sulfuric acid from elemental sulfur is to produce it by burning H2S. Reaction R-6.27 would be replaced by: H2 SðgÞ þ 2O2 ðgÞ/SO2 ðgÞ þ H2 OðgÞ

R-6.32

The advantage to this approach to sulfuric acid production is that it would eliminate the need for Claus and SCOT plants, thereby potentially reducing costs. Very few sources of H2S, except for a few large sour natural gas fields, produce sufficient H2S to supply a phosphoric acid plant. Consequently, burning H2S to make sulfuric acid makes sense only if there is a local small volume demand for the acid.

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CO-SEQUESTRATION OF CO2 AND H2S Most current plans for coal gasification include elemental sulfur production. In the economic analysis, the sale of sulfur is often included as an economic benefit; and this benefit is included as a substantial advantage for the gasification of high sulfur coals. The problem with this analysis is that sulfur prices are very dependent on geographic location. For example, large quantities of sulfur are produced during the purification of sour natural gas in the Rocky Mountain region of the USA. The US phosphate industry, centered in Tampa, Florida, is the primary destination for elemental sulfur in the USA. The delivered price for sulfur in Tampa roughly equals the cost of shipping sulfur from the Rocky Mountain producers to Florida; so the sulfur producers receive essentially nothing for their product. The sour natural gas producers continue to ship sulfur to avoid the responsibility of maintaining sulfur piles. In Alberta, Canada, large quantities of sulfur are produced as a byproduct of sour natural gas and oil sands processing. These producers ship sulfur via rail to the port of Vancouver for sales to Pacific Rim countries. When delivered, prices aren’t sufficient to cover shipping charges, and large quantities of sulfur pile up near producers’ facilities. With the increasing use of coal gasification, sulfur production is expected to rise. Since the demand for sulfur is not expected to rise substantially, sulfur prices are likely to fall; and fewer sulfur producers will be able to sell sulfur at prices sufficient to cover shipping costs. An alternative approach is to co-sequester H2S with CO2. In this case, sulfur is not viewed as a valuable product, but rather a waste stream that requires disposal. The Great Plains Synfuels plant in North Dakota pipelines a mixture of CO2 and H2S to the Weyburn enhanced oil project in Saskatchewan.40 Hydrogen sulfide in this gas enhances the oil solvency of the gas mixture, thereby improving oil recovery. Hydrogen sulfide, however is dissolved in the produced oil and must be removed in the refinery. It consequently lowers the value of the produced oil. Sequestration of CO2 in deep saline aquifers eliminates this problem. The co-sequestration of CO2 and H2S simplifies a gasification plant by eliminating the need for selective acid gas removal, and by eliminating the need for a Claus plant and a SCOT plant. This reduces the capital and operating cost for the plant, but hydrogen sulfide in the CO2 stream does impose additional demands. Compressor, pipeline, and pump materials must be upgraded to withstand the corrosive properties of H2S. There are also substantial safety issues. Carbon dioxide gas can be deadly, but only at high concentrations. Hydrogen sulfide, on the other hand, is a neurotoxin that is dangerous at very low concentrations. A sequestration pipeline that runs several kilometers from the producer to the injection site needs numerous safeguards to protect the public. Enhanced oil recovery is an attractive form of CO2 sequestration because the cost of sequestration is offset by increased oil production. Other forms of sequestration also need

Sulfur Recovery

to be considered as there is not sufficient capacity in old oil fields to accept all of the CO2 that needs to be sequestered; and enhanced oil recovery is only feasible when the CO2 source is near an oil field. The oil fields can be roughly divided into sour (high sulfur) and sweet (low sulfur) fields. Sour oil producers are already equipped to handle H2S, so they are likely to welcome CO2/H2S mixtures for enhanced oil recovery. A sweet oil producer is less likely to use a CO2/H2S mixture, because this will require a large investment in H2Sresistant injection and production facilities. Already, a number of sour natural gas producers re-inject (sequester) H2S rather than produce elemental sulfur. This practice will probably become more common with increased sulfur production and falling sulfur prices.

REFERENCES 1. U.S. Energy Information Administration. Coal production and number of mines by State and mine type. <www.eia.doe.gov/cneaf/coal/page/acr/table1.html; 2008>. 2. Overview: The Clean Air Act Amendments of 1990. <www.epa.gov/oar/caa/overview.txt>. 3. Kohl A, Reisenfeld F. Gas Purification. 4th ed. Gulf Publishing; 1985. 4. Undengaard NR, Berzins V. Catalytic conversion of COS for gas clean-up. In: Newman SA, ed. Acid and Sour Gas Treating Processes. Gulf Publishing; 1985. p. 603-616. 5. Carnell PJH. Feedstock purification. In: Twigg MV, ed. Catalyst Handbook. 2nd ed. Wolfe Publishing; 1989. p. 191-224. 6. Leonard JP, Haritatos NJ, Law DV. The Chevron WWT processdAn economical way to treat sour water from refineries and synthetic fuel plants. In: Newman SA, ed. Acid and Sour Gas Treating Processes. Gulf Publishing; 1985. p. 734-752. 7. Ranke G, Mohr VK. The Rectisol wash: New developments in acid gas removal from synthesis gas. In: Newman SA, ed. Acid and Sour Gas Treating Processes. Gulf Publishing; 1985. p. 80-111. 8. Lurgi. The Rectisol process: Lurgi’s leading technology for purification and conditioning of synthesis gas. <www.lurgi.com/website/fileadmin/user_upload/1_PDF/1_Broshures_Flyer/englisch/0308e_ Rectisol.pdf>. 9. Breckenridge W, Holiday A, Ong JOY, et al. Use of SELEXOL process in coke gasification to ammonia project, Laurance Reid Gas Conditioning Conference. Norman, OK: University of Oklahoma. <www.uop. com/objects/92selexcokegasifamm.pdf;; 2000>. 10. Clark M. UOP Selexol technology applications for CO2 capture, 3rd Annual Wyoming CO2 Conference, Casper, WY; 2009. <eori.uwyo.edu/downloads/CO2_Conf_2009/Presentation PDF/ Mike Clark REVISED Selexol_Wyoming_Conference.pdf>. 11. Palla R. Meeting staged CO2 capture requirements with the UOP SELEXOL process. Gasification Technologies Council Annual Conference; 1999. 12. Maddox RN, Bhairi A, Mains GJ, et al. Equilibrium between CO2/H2S and ethanolamine solutions. In: Newman SA, ed. Acid and Sour Gas Treating Processes. Gulf Publishing; 1985. p. 212-234. 13. Polasek J, Bullen J. Process considerations in selecting amines. In: Newman SA, ed. Acid and Sour Gas Treating Processes. Gulf Publishing; 1985. p. 190-211. 14. Sartori G, Savage DW. Sterically hindered amines for CO2 removal from gases. Ind. Engr. Chem. Fundam. 1983;vol. 22:239-249. 15. Goldstein AM, Edelman AM, Ruziska PA. FLEXSORB gas treating: New amine technologies. In: Newman SA, ed. Acid and Sour Gas Treating Processes. Gulf Publishing; 1985. p. 319-341. 16. Benson HE, Field JH, Hayes WP. Improved process for CO2 absorption uses hot carbonate solutions. Chem Engr Prog 1956;52:433-438.

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17 UOP. Gas processing: Benfield process. <www.uop.com/objects/99 Benfield.pdf>. 18. Gangriwala HA, Chao I-M. The CATACARB process for acid gas removal. In: Newman SA, ed. Acid and Sour Gas Treating Processes. Gulf Publishing; 1985. p. 370-387. 19. Danckwerts PV, Sharma MM. The absorption of carbon dioxide into solutions of alkalis and amines (with some notes on hydrogen sulphide and carbonyl sulphide). Chem Engr CE244; 1966. 20. Tseng PC, Ho WS, Savage DW. Carbon dioxide absorption into promoted carbonate solutions. AIChEJ 1988;vol. 34:922-931. 21. Sartori G, Savage DW. Sterically hindered amines for CO2 removal from gases. Ind Eng Chem Fundam 1983;22:239-249. 22. Cullinane JT, Rochelle GT. Carbon dioxide absorption with aqueous potassium carbonate promoted by piperazine. Chem Engr Sci 2004;vol. 59:3619-3630. 23. Sharma MM, Danckwerts PV. Catalysis by Bro¨nsted bases of the reaction between CO2 and water. Trans Faraday Soc 1963;59:386-395. 24. Dennard AE, Williams RJP. The catalysis of the reaction between carbon dioxide and water. J Chem Soc, Sect A. 1966:pp.812-816. 25. Astarita G, Savage DW, Longo JM. Promotion of CO2 mass transfer in carbonate solutions. Chem Engr Sci 1981;36:581-588. 26. Cents AHG, F Brilman DW, Versteeg GF. CO2 absorption in carbonate/bicarbonate solutions: The Danckwerts-Criterion revisited. Chem Engr Sci 2005;60:5830-5835. 27. Powerspan. <www.powerspan.com>. 28. Pennline H, Richards G, Plasynski SI. Ammonia-based process for multicomponent removal from flue gas. <www.netl.doe.gov/publications/factsheets/rd/R&D043.pdf>. 29. Yi C-K, Jo S-H, Jin G-T, et al. Continuous operation of spray-dried zinc based sorbent in a hot gas desulfurization process consisting of a transport desulfurizer and a fluidized regenerator. 5th Int. Symp. Gas Cleaning at High Temp. <www.netl.doe.gov/publications/proceedings/02/GasCleaning/gas_ clean02.html#Session 6; Sept. 17e22, 2002>. 30. Gupta R, Turk B, Lesemann M. RTI/Eastman warm syngas clean-up technology: Integration with carbon capture. Gasification Technologies Annual Conference; 2009. 31. Kobayashi M, Shirai H, Nunokawa M. Stability of sulfur capacity attributed to zinc sulfidation on sorbent containing zinc ferrite-silica composite powder in pressurized coal gas. 5th Int Symp Gas Cleaning at High Temp. . 32. Shirai H, Kobayashi M, Nunokawa M, et al. Regeneration and durability of advance zinc ferrite sorbent for hot coal gas desulfurization. 5th Int Symp Gas Cleaning at High Temp, <www.netl.doe.gov/ publications/proceedings/02/GasCleaning/gas_clean02.html#Session 6>; Sept. 17e22, 2002. 33. Cheah S, Carpenter DL, Magrini-Bair KA. Review of mid- to high-temperature sulfur sorbents for desulfurization of biomass- and coal-derived syngas. Energy & Fuels; 2009. doi:10.1021/ef900714q. 34. Lew S, Jothimurugesan K, Flytzani-Stephanopoulos M. High-temperature H2S removal from fuel gases by regenerable zinc oxide-titanium dioxide sorbents. Ind Eng Chem Res 1989;vol. 28:535-541. 35. Wakker JP, Gerritsen AW, Moulijn JA. High temperature H2S and COS removal with MnO and FeO on g-Al2O3 acceptors. Ind Eng Chem Res 1993;32:139-149. 36. Fenouil LA, Lynn S. Study of calcium-based sorbents for high-temperature H2S removal. 1. Kinetics of H2S sorption by uncalcined limestone. Ind Eng Chem Res 1995;34:2324-2333. 37. Fenouil LA, Lynn S. Study of calcium-based sorbents for high-temperature H2S removal. 2. Kinetics of H2S sorption by calcined limestone. Ind Eng Chem Res 1995;34:2334-2342. 38. Capone M. In: Kroschwitz JI, Howe-Grant M, Humphreys L, et al., eds. Sulfur removal and recovery, Kirk-Othmer Encyclopedia of Chemical Technology. 4th ed. vol. 23. John Wiley and Sons; 1997. p. 432452. 39. Ober JA. Materials flow of sulfur. US Geological Survey Open File Report; 2002:02-298. 40. Plasynski S, Brickett LA, Preston CK. Weyburn carbon dioxide sequestration project. <www.netl.doe. gov/publications/factsheets/project/Proj282.pdf; 2008>.

CHAPTER

7

Hydrogen Production and Integrated Gasiﬁcation Combined Cycle (IGCC) Contents Need for Increasing H2 Content Water Gas Shift in the Catalytic Temperature Range Hydrogen for Ammonia Synthesis Iron-Based HT Shift Catalyst LT Shift Catalyst Sour Gas Shift Steam-Iron Process Hydrogen for Ammonia Synthesis: Removal of Residual Impurities Dehydration Hydrogen for Proton Exchange Membrane Fuel Cells Hydrogen for Petroleum Reﬁning Combined cycle plants for power production, NGCC, and IGCC Natural Gas Combined Cycle (NGCC) Integrated Gasiﬁcation Combined Cycle (IGCC) Combining IGCC and Oxy-Combustion Methanol, SNG, and Fischer-Tropsch Synthesis References

137 138 139 140 142 142 143 145 146 146 148 149 151 152 154 155 155

NEED FOR INCREASING H2 CONTENT Most applications of coal-derived syngas require either a nearly pure hydrogen stream, or a H2/CO mixture with a greater hydrogen content than that provided by the gasifier. In all these cases, the hydrogen concentration needs to be increased, and impurity gases must be removed. Table 7.1 shows an estimate of the H2/CO syngas ratios for several gasifiers. The temperatures shown are for the syngas prior to a quench for gasifiers in which a quench is an integral part of the reactor. The water gas shift reaction was described in Chapter 3. COðgÞ þ H2 OðgÞ4CO2 ðgÞ þ H2 ðgÞDHrxn ¼ 41:21 kJ=gmole

R-7.1

The equilibrium constant, KWG, may be described in terms of gas partial pressures, Pi, or gas molar fractions, yi: KWG ¼

yCO2 yH2 PCO2 PH2 ¼ PCO PH2 O yCO yH2 O

Coal Gasification and Its Applications. ISBN B978-0-8155-2049-8.10007-5, doi:10.1016/B978-0-8155-2049-8.10007-5

Eqn. 7.1

Ó 2011 Elsevier Inc. All rights reserved.

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Hydrogen Production and Integrated Gasiﬁcation Combined Cycle (IGCC)

Table 7.1 Syngas H2/CO and CO2/CO ratios for several gasiﬁers. Gasiﬁer GE1 Conoco-Philips1 Shell1 Entrained Entrained Entrained Type ﬂow ﬂow ﬂow

Transport2 Fluidized bed

BGL3 Moving bed

Temperature, C H2/CO CO2/CO H2O/CO

907 1.65 1.67 3.59

537 0.41 0.12 0.003

1,316 0.97 0.44 0.42

1,010 0.71 0.38 0.32

1,427 0.51 0.04 0.06

The equilibrium constant falls as the temperature increases. In the gasifier, temperatures are high enough for the water gas shift reaction to proceed to near equilibrium values without a catalyst. As can be seen in Table 7.1, the high temperatures of entrained flow gasifiers generally produce H2/CO ratios less than 1. The Transport gasifier, on the other hand, due to a relatively low operating temperature and a high steam feed rate, produces a relatively high H2/CO ratio. The BGL gasifier has an even lower operating temperature, but the steam rate in this example is very low, producing a low H2/CO ratio. Applications that require nearly pure H2 include ammonia synthesis and PEM fuel cells. Somewhat less rigorous standards apply to hydrogen produced for petroleum hydrogenation and power production via integrated gasification combined cycle (IGCC) plants. Methanol, substitute natural gas (SNG) and Fischer-Tropsch synthesis require a H2/CO blend.

WATER GAS SHIFT IN THE CATALYTIC TEMPERATURE RANGE To increase the H2/CO ratio, syngas is fed to a catalytic water gas shift reactor. Figure 3.2 shows the equilibrium constant as a function of temperature for the 800 to 1,500 C range, which corresponds to the typical range of gasification temperatures. To further increase the H2/CO ratio, catalytic water gas shift reactors operate at lower temperatures. Figure 7.1 shows the equilibrium constant as a function of temperature for the 200 to 600 C range. Eqn. 3.2 is an empirical fit for the equilibrium versus temperature results for the 800 to 1,500 C range. The corresponding fit (same equation form, different constants) for the 200 to 600 C range is: lnðKWG Þ ¼ 4:151 ½lnðT Þ2 60:913 ½lnðT Þ þ 223:14 Where T ¼ temperature, K.

This equation fits the calculated equilibrium constants to within 0.3%.

Eqn. 7.2

Equilibrium Constant

Hydrogen Production and Integrated Gasiﬁcation Combined Cycle (IGCC)

100

10

1

200

250

300

350

400

450

500

550

600

Temperature, C

Figure 7.1 Water gas shift reaction equilibrium constant as a function of temperature for the 200 to 600 C range at 2 MPa. See Figure 3.2 for the 800 to 1,500 C range.

HYDROGEN FOR AMMONIA SYNTHESIS Conversion of syngas to hydrogen for ammonia synthesis is one of the oldest applications of the water gas shift reaction, and remains as one of the applications with the most stringent product specifications. Consequently, a description of this process serves as a good base case for comparison to other applications Figure 7.2 shows a typical flowsheet for conversion of syngas to hydrogen. The water gas shift reaction occurs in two reactors, a high temperature (HT) shift reactor which accomplishes the bulk of the reaction; and a low temperature (LT) shift reactor that pushes the water gas shift reaction further to the right. The traditional HT shift catalyst can tolerate low levels of sulfur; but the LT shift catalyst is poisoned by sulfur, so sulfur is removed from the syngas prior to the shift reactors using an acid gas removal (AGR) process. In natural gas derived syngas, the sulfur levels are sufficiently low enough that the ZnO bed may be operated without the first AGR process. The higher H2S levels in coalderived syngas make the first AGR process necessary. If the first AGR process also removes CO2, then this pushes the equilibrium toward increased H2 production. The product of the LT shift reactor is sent to a second AGR process to remove CO2. As will be described later, the H2 stream requires further processing before it may be sent to ammonia synthesis. Different catalysts are used for the HT and LT shift reactors. The LT catalyst is much more active; but it is also easily sintered at elevated temperatures, including the temperature rise that occurs due to the exothermic heat of reaction. Consequently, the bulk of the reaction is accomplished using a less active, but more thermally robust catalyst in the HT shift reactor.

139

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Hydrogen Production and Integrated Gasiﬁcation Combined Cycle (IGCC)

H2S, CO2 syngas

AGR process

H2, CO

ZnO bed steam

370 °C high temperature shift

low temperature shift

CO2

AGR process

440 °C

H2

220 °C 240 °C

Figure 7.2 Conversion of coal-derived syngas to hydrogen for ammonia synthesis.

IRON-BASED HT SHIFT CATALYST The traditional HT catalyst4 is supplied as a blend of 90e95% Fe2O3 (hematite) and 5e10% Cr2O3. This catalyst is activated in situ by syngas, which reduces hematite to magnetite and chromium VI to chromium III. Once activated, the catalyst surface area is typically 60 to 120 m2/g. ¼ 7:1 KJ=mole 3Fe2 O3 ðsÞ þ H2 ðgÞ42Fe3 O4 ðsÞ þ H2 OðgÞ DHrxm

R-7.2

¼ 48:1 KJ=mole DHrxm

R-7.3

3Fe2 OðsÞ þ COðgÞ42Fe3 O4 ðsÞ þ CO2 ðgÞ

2CrO3 ðsÞ þ 3H2 ðgÞ/Cr2 O3 ðsÞ þ 3H2 OðgÞ

R-7.4

2CrO3 ðsÞ þ 3COðgÞ/Cr2 O3 ðsÞ þ 3CO2 ðgÞ

R-7.5

To prevent over-reduction of the catalyst (conversion of Fe3O4 to FeO), a minimum H2O/H2 ratio is required.4 This ratio increases from about 0.1 at 360 C to about 0.9 at 550 C.

Hydrogen Production and Integrated Gasiﬁcation Combined Cycle (IGCC)

The catalyst is a solid solution of Cr2O3 in Fe3O4. The role of Cr2O3 appears to be the prevention of catalyst sintering, which leads to a loss of surface area.5 Podolski and Kim6 fit experimental rate data to Langmuir-Hinshelwood and power law kinetic models. Both models gave comparable data fits. The power law model, the simpler of the two, is: 118; 943J=gmole ðPCO Þ0:81 6 r ¼ 3X10 exp Eqn. 7.3 RT ðPH2 O Þ0:024 ðPCO2 Þ0:16 ðPH2 Þ0:044 Where r ¼ conversion rate of CO, gmole/(g catalyst)-min

R ¼ gas constant, 8.31447 J/gmole-K T ¼ temperature , Kelvin Pi ¼ partial pressure of gas i, atmospheres. The catalyst used by Podolski and Kim had a 18 m2/g surface area, versus 60 to 120 m2/g in commercial catalysts. As a first approximation, the pre-exponential constant should be a linear function of the surface area. For example, a commercial catalyst with a 90 m2/g surface area, five times greater than catalyst used by Podolski and Kim, therefore it should have a pre-exponential constant five times greater, 1.5 107, versus the 3 106 shown in Eqn. 7.3 The effectiveness factor, h, is defined as: robs ¼ hrsl

Eqn. 7.4

Where robs ¼ the observed reaction rate

rsl ¼ the surface limited reaction rate that would be observed with no mass transfer restriction. For the water gas shift reaction in the HT shift reactor, rsl in Eqn. 7.4 is the same as r in Eqn. 7.3. Lloyd et al.4 reported effectiveness factors ranging from near unity at 310 C to 0.24 to 0.45 at 450 C for two commercial iron based HT shift catalysts. Hydrogen sulfide temporarily poisons the iron-based shift catalyst via the reversible reaction of H2S with magnetite to form iron sulfide: Fe3 O4 ðsÞ þ 3H2 SðgÞ þ H2 ðgÞ43FeSðsÞ þ 4H2 OðgÞ

¼ 89:6 KJ=mole DHrxn R-7.6

Bohlbro7 measured the effects of 0 to 2,000 ppm H2S on rate of the water gas shift reaction. The HT shift reactor for a 1,000 ton/day ammonia plant, using natural gas as the feedstock, typically8 has a volume of about 57 m3. Assuming complete conversion of CO to H2, and complete conversion of H2 to NH3, this is the equivalent of 0.029 m3 of reactor volume per kg-mole/h of H2 plus CO. With a typical4 bulk catalyst density of 1.15 kg/liter, this is equivalent to 33 kg of catalyst per kg-mole/h of H2 plus CO. This is sufficient to achieve a near-equilibrium composition for the water gas shift reaction.

141

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Hydrogen Production and Integrated Gasiﬁcation Combined Cycle (IGCC)

LT SHIFT CATALYST The traditional LT shift catalyst4 consists of roughly equal portions of CuO, ZnO, and Al2O3. The catalyst is activated by reducing CuO to metallic copper: CuOðsÞ þ H2 ðgÞ/CuðsÞ þ H2 OðgÞ

R-7.7

Metallic copper is the active catalytic species. The remaining components are present to retard catalyst sintering. The catalyst is poisoned by sulfur and halides. The LT shift catalyst is especially prone to poisoning; in part, due to the low operating temperatures, which are insufficient to desorb some poisons. Even with poison-free gas, the catalyst deactivates due to sintering. Rapid sintering occurs above 260 C, and slower sintering occurs at normal operating temperatures. Minimum temperatures for gas entering the LT shift reactor are limited by the dew point of water in the gas. The gas temperature rises as it passes through the reactor due to the exothermic heat of reaction; and must be well below 260 C at the exit to avoid sintering. This limits the LT shift reactor to a narrow band of operating temperatures and a low CO conversion across the reactor. Consequently, most of the water gas shift reaction must occur in the HT shift reactor. Pearce et al.8 stated that a typical LT shift reactor volume for a 1,000 ton/day ammonia plant is 61 m3, which is a little greater than the typical volume for a HT shift reactor. The LT shift reactor contains about three times the catalyst needed, based on the activity of fresh catalyst, to convert the feed gas to near equilibrium levels. This volume of catalyst allows the time between catalyst replacement to be extended. Early in the bed life, most of the reaction occurs near the reactor inlet. The catalyst near the inlet deactivates first, and the most reactive portion of the bed moves toward the outlet.

SOUR GAS SHIFT Sulfur tolerant water gas shift, also known as sour gas shift: catalysts are an attractive option for coal-derived synthesis gas. These catalysts are typically bimetallic, sulfided Co/Mo on an alumina support; essentially the same catalyst long used to desulfurize petroleum products. Newsome5 describes a series of laboratory experiments in which a commercial petroleum desulfurization catalyst is promoted by impregnating the catalyst with an alkali salt. The resulting catalyst was more active than the copper-based LT catalyst, and could be used at temperatures as high as 400 C. Sour gas shift catalysts do not require the first AGR removal process and the ZnO bed shown in Figure 7.2. A single AGR process after the sour gas shift is used to remove H2S and CO2. This can be a considerable cost savings. Another benefit is that COS hydrolysis (reaction R-6.3) occurs within the sour gas shift reactor, so a separate COS hydrolysis reactor is not required.

Hydrogen Production and Integrated Gasiﬁcation Combined Cycle (IGCC)

Haldor Topsoe, Inc., a commercial sour gas shift catalyst manufacturer, suggests9 two flow sheets for a sour gas shift operation. The first flow sheet uses a sour HT shift in the initial stage and is following by a sour LT shift converter in the secondary stage. Both reactors appear to contain the same catalyst. The HT conditions favor a higher rate of reaction; while the LT conditions push the water gas shift equilibrium further toward the right. The second flowsheet uses a single shift reactor, and is for applications that require a lower H2/CO product ratio.

STEAM-IRON PROCESS The steam-iron process is an alternative to the water gas shift reaction for producing H2 from syngas. In newer literature, it is also known as the syngas chemical looping (SCL) process. The first commercial plant10 was built in 1904, but there does not appear to be any recent commercial application. Recent preliminary process and economic studies11e13 show that the steam-iron process may be more attractive than the conventional water gas shift/AGR combination. The process concept revolves around the reversible oxidation of a metal. Several metals have been considered, but iron is preferred.14 The process is shown in Figure 7.3. Syngas is desulfurized, and then hot (750 to 900 C) syngas, at about 3 MPa, contacts a countercurrent stream of solid Fe2O3 in the reducer.15 This reactor converts the gas stream to CO2 and H2O, and the solid stream to a mixture of Fe and FeO. 3Fe2 O3 ðsÞ þ H2 ðgÞ42Fe3 O4 ðsÞ þ H2 OðgÞ 3Fe2 OðsÞ þ COðgÞ42Fe3 O4 ðsÞ þ CO2 ðgÞ

¼ 7:1 KJ=mole DHrxn

R-7.8

¼ 48:1 KJ=mole DHrxn

R-7.9

Fe3 O4 ðsÞ þ H2 ðgÞ43FeOðsÞ þ H2 OðgÞ

¼ 62:9 KJ=mole DHrxn

R-7.10

Fe3 O4 ðsÞ þ COðgÞ43FeOðsÞ þ COðgÞ

¼ 21:8 KJ=mole DHrxn

R-7.11

¼ 30:2 KJ=mole DHrxn

R-7.12

¼ 10:9 KJ=mole FeOðsÞ þ COðgÞ4FeðsÞ þ CO2 ðgÞ DHrxn

R-7.13

FeOðsÞ þ H2 ðgÞ4FeðsÞ þ H2 OðgÞ

Hot gas leaving the reducer is used to generate steam. If this gas is sufficiently cooled, most of the water will condense, which leaves a CO2 stream that can be sequestered. The Fe/FeO solid mixture then flows to the oxidizer, where the solids contact steam at about 500 to 750 C and 3 MPa. The steam then reacts with the solid to form H2, the primary product, and Fe3O4 (magnetite). The magnetite is entrained in a stream of compressed air. This gas/solid mixture flows to the combustor, where magnetite is oxidized to Fe2O3 (hematite) at about 950 to

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Hydrogen Production and Integrated Gasiﬁcation Combined Cycle (IGCC)

water

steam

CO2

CO2, H2O water

Turbine

Reducer

H2, CO

Fe, FeO

Fe2O3 steam Combustor water

H2 O2-depleted air Oxidizer

steam Fe3O4 compressor air

Figure 7.3 The steam-iron process for the production of H2 from syngas.

1,150 C and 3.2 MPa. A hydrocyclone is used to separate the solid hematite and the O2-depleted air. Hematite flows to the reducer to complete the solids loop. The O2-depleted air is routed through a turbine to recover air compression energy; and then used to generate steam to recover thermal energy. Gasior et al.16 described a series of pilot plant experiments. Only about 65% of the syngas was converted in the reducer, which would leave substantial quantities of unreacted H2 and CO in the CO2 product stream. They did not use a combustor step, so the reducer was fed magnetite, rather than hematite, and this may account for the low syngas conversion rate. Solids conversion in the reducer was 25 to 30%. Hydrogen purity from the oxidizer was 91 to 98%. Removal of CO2, the primary impurity, would increase the H2 purity to over 95%. Li et al.15 describe the synthesis of a high surface area pellet containing about 60% Fe2O3 and 40% Al2O3 using sol gel techniques. In fixed bed tests at 830 C, essentially all

Hydrogen Production and Integrated Gasiﬁcation Combined Cycle (IGCC)

of the H2 and CO was converted to H2O and CO2. This occured until breakthrough was observed. They showed that pure iron oxides tended to deactivate with repeated cycling, while their iron/aluminum oxide solids maintained activity. Hydrogen produced during the oxidation step had a purity greater than 99.6%. The steam-iron process requires a sulfur-free syngas feed stream. Wakker et al.17 used FeO to remove H2S from syngas. This work suggests that if the syngas feed contained H2S, then the sulfur would be trapped as FeS in the reducer. It would then be converted to FeO and H2S in the reducer stage, thereby contaminating the H2 product with H2S. In the process and economic study by Tomlinson et al.,12 sour syngas is cooled and fed to an MDEA-based AGR process to remove sulfur. The sulfur-free syngas, at 66 C, is re-heated to 533 C and fed to the steam-iron process. This process would be more efficient if a high temperature sulfur removal process were used.

HYDROGEN FOR AMMONIA SYNTHESIS: REMOVAL OF RESIDUAL IMPURITIES Oxygen compounds, including CO, CO2, and H2O, are temporary poisons for the ammonia synthesis catalyst, meaning that they depress catalytic activity; but the activity is restored once these gases are eliminated.18 Syngas leaving the LT shift reactor still contains 0.1 to 0.3% CO, and syngas leaving the AGR contains considerable water vapor and traces of CO2. These impurities must be removed prior to ammonia synthesis. One option to purify the hydrogen is to use a pressure swing adsorption process. The adsorbent is a molecular sieve with pores just large enough to adsorb H2. The other gas molecules, which have larger molecular diameters, pass through the adsorbent. Feed to the bed is shut off, and the bed pressure is reduced to desorb H2. By using multiple beds, this cyclic process can continuously purify hydrogen. A more common approach in ammonia plants is methanation. Carbon monoxide and carbon dioxide are converted to methane and water; ¼ 206:2 KJ=mole COðgÞ þ 3H2 ðgÞ/CH4 ðgÞ þ H2 OðgÞ DHrxn

CO2 ðgÞ þ 4H2 ðgÞ/CH4 ðgÞ þ 2H2 OðgÞ

¼ 165:1 KJ=mole DHrxn

R-7.14 R-7.15

Note that R-7.14 is the reverse of the methane steam reforming reaction, R-11.2. Low temperatures favor methane formation. Typically,8 a 320 C methanation temperature is used to get reasonably fast reaction rates and low equilibrium CO level. A nickel catalyst is typically used for methanation. This catalyst is poisoned by sulfur, but the feed to methanation typically contains little or no sulfur.

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Hydrogen Production and Integrated Gasiﬁcation Combined Cycle (IGCC)

dry gas

dry glycol

cooler

absorption column

steam

moist gas

boiler flash tank wet glycol

Figure 7.4 Glycol dehydration process for the removal of water vapor from a gas stream.

DEHYDRATION Methane in the ammonia synthesis reactor behaves as an inert, but water is a catalyst poison and must be removed. Much of the water can be removed by simply cooling the gas and condensing water. Further water reduction requires a dehydration process. Glycols are commonly used to dehydrate gas streams. The glycol is usually either ethylene glycol, HO-CH2CH2-OH, also known as monoethylene glycol, MEG, or EG, or triethylene glycol, HO-(CH2CH2-O)3H, also known as TEG. Glycols are quite similar to Selexol solvent, will replace the terminal hydrogen with a methyl group. Ethylene glycol has a greater water affinity, but triethylene glycol has a lower vapor pressure, which reduces solvent loss. Figure 7.4 shows a glycol dehydration process. Note the similarity between this process and a typical acid gas removal process. Solid adsorbents, such as silica gel and alumina, are sometimes used instead of glycols to dehydrate gas streams.

HYDROGEN FOR PROTON EXCHANGE MEMBRANE FUEL CELLS Much has been written about the concept of a hydrogen economy; in which H2 is the fuel of choice. More specifically, the concept of using H2 as a transportation fuel has received a great deal of attention. Some H2-as-transportation-fuel concepts involve burning H2 in an internal combustion engine; but the usual approach is to feed H2 to a proton exchange membrane (PEM) fuel cell, and use the electricity produced by the fuel cell to power an electric motor.

Hydrogen Production and Integrated Gasiﬁcation Combined Cycle (IGCC)

Proponents claim that since H2 burns to form water, has no harmful emissions. This is not strictly true. If H2 is burned in air, a small quantity of NOx emissions will form. In a PEM fuel cell, NOx formation is insignificant; but the exhaust may contain harmful impurities, such as methane, a potent greenhouse gas. A more significant source of harmful emissions is the process used to produce H2. For example, large quantities of H2 can be made by electrolyzing water using off-peak electric power from existing power plants. If the electric power comes from a coal-fired power plant; then most of the emissions from H2 use are due to the increased combustion of coal at the power plant. Still, it is usually more practical to control emissions from a large point source, such as a power plant, than it is from a multitude of small sources, such as vehicles and home furnaces. Hydrogen is a low emissions technology only if the H2 is produced using low emissions techniques. This is a book about coal gasification, so the focus will be on H2 produced from coalderived syngas. Conventional PEM fuel cells use platinum electrodes and a polymeric electrolyte. The electrodes are poisoned by CO, which is a common component of syngas. Since CO also poisons ammonia synthesis gas; the techniques used to produce H2 for ammonia synthesis are also applicable to H2 for PEM fuel cells. The major difference is that PEM fuel cells require a humidified H2 feed; while the H2 used for ammonia synthesis must be dry. Above 150 C, CO desorbs from platinum surfaces, so CO is no longer a poison. However, the polymeric electrolyte cannot withstand these elevated temperatures. At the time this book was written, two major technical challenges blocked commercialization of PEM fuel cell vehicles. The first is that worldwide production of platinum is not sufficient to produce a large number of PEM fuel cell vehicles. Advances in fuel cell technology may one day overcome this obstacle. The second major challenge is the need for an economical, safe, and practical means of storing sufficient H2. Currently storage of H2 in vehicles is not enough to give them a reasonable driving range between refueling stops. This challenge is explored in Chapter 8. Ideally, an H2-powered vehicle would have a driving range comparable to what is normally expected from a gasoline or diesel powered vehicle, around 500 to 750 km per fueling. To be commercially viable, the range for an H2-powered vehicle needs to be at least as great as a battery-powered vehicle, now about 150 km. Gasoline and diesel are convenient transportation fuels because they are liquids at ambient temperature and pressure, and because of their high energy density (available energy per unit mass or volume). A typical heat of combustion for gasoline is 44 MJ/kg (31 MJ/L).19 Diesel has a slightly higher heat of combustion on a mass basis, and a substantially higher heat of combustion on a volume basis due to its higher density. A typical C-class automobile has a 45 liter gasoline tank, giving it an energy storage capacity of about 1,400 MJ and a 33 kg fuel weight. Hydrogen has a much higher energy on a mass basis, 120 MJ/kg, but it is much more difficult to store. The US Department of Energy set a 5 wt.% hydrogen goal for

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Hydrogen Production and Integrated Gasiﬁcation Combined Cycle (IGCC)

hydrogen storage systems, which would give a 6 MJ/kg energy density. To achieve 1,400 MJ energy storage, the fuel system would weigh 233 kg. The extra 200 kg would be equivalent to hauling around an extra 2-3 people everywhere you go. Of course, PEM fuel cells are substantially more efficient than gasoline engines, so the weight penalty for an equivalent range will not be so high. Still, as we will see in Chapter 8, achieving the 5 wt.% hydrogen storage goal is not a trivial task.

HYDROGEN FOR PETROLEUM REFINING Catalytic hydrogenation of petroleum is widely used in petroleum refining to remove sulfur, nitrogen, and oxygen from feedstocks, to saturate olefins and aromatics; and to reduce the molecular weight of heavy feedstocks, hydrocracking. Most refineries employ a naphtha reforming process. This process dehydrogenates alkanes to produce aromatic hydrocarbons in order to increase the octane number of the finished gasoline product. If a light, sweet crude oil is refined, then the hydrogen produced as a byproduct of naphtha reforming may be sufficient to meet the refinery’s hydrogen needs. The general trend in the petroleum industry has been toward heavier, sour crude oils. A notable component of this trend is the growth of the Canadian oil sands industry. When a heavy oil is refined, hydrogen from the naphtha reformer may not be sufficient, and a separate hydrogen production plant is required. Heavy, high molecular weight, petroleum cuts generally have a lower than desired hydrogen/carbon ratio. To produce products with the desired hydrogen/carbon ratio, a refiner will generally employ either a carbon-out or a hydrogen-in refining strategy. Carbon-out strategies tend to have lower refining costs, but also give lower product yields. Crude oil is distilled in an atmospheric pressure distillation column to produce fuel gas, naphtha (straight run gasoline), jet fuel (kerosene), and diesel. The bottoms product of the atmospheric crude tower, known as atmospheric residuum (resid) or topped crude, is sent to a vacuum distillation tower to produce vacuum distillates and vacuum residuum. Vacuum distillates have normal (atmospheric pressure equivalent) boiling points in the 400 to 570 C range, and carbon numbers in the C20 to C40 range. Vacuum resids have normal boiling points in the 550 Cþ range, carbon numbers in the C40þ range, and may contain a small quantity of solids. A common, carbon-out process to refine vacuum distillates is fluidized catalytic cracking (FCC), a catalyzed pyrolysis process. The feedstock pyrolyzes on the catalyst surface to produce gas, liquids, and solid carbon. The carbon will deposit itself on the catalyst surface. The catalyst increases the liquid yield compared to a non-catalyzed pyrolysis process. Carbon-loaded catalyst is sent to a separate vessel where the carbon is burned off. This regenerates the catalyst and provides heat for the endothermic pyrolysis reactions. FCC is rarely used with vacuum resid feeds because of rapid, irreversible poisoning of the catalyst. Instead a coking, which is a non-catalytic pyrolysis, process is used. This

Hydrogen Production and Integrated Gasiﬁcation Combined Cycle (IGCC)

process produces a solid petroleum coke, or petcoke, byproduct. Petcoke can be burned in a coal-fired power plant. But, despite its high heat of combustion, petcoke is not a desirable fuel due to its high sulfur content, low volatile content, and low reactivity. Because of its abundance and low value, petcoke is often mentioned as a potential gasification feedstock, which makes a combined carbon-out, hydrogen-in refining approach possible. Arienti20 presented a design study in which petcoke is gasified to provide H2, steam, and electricity for a refinery. Hydrogen was produced by concentrating H2 in syngas using a membrane, and then a pressure wing adsorption unit was used to purify the H2. In the base case, there was no water gas shift. The CO rich stream was burned to produce electricity. Atmospheric and vacuum distillates are normally hydrogenated in a trickle bed reactor. Oil and a large stoichiometric excess of hydrogen are fed to the top of a reactor packed with a solid bed of hydrogenation catalyst. The oil then trickles through the catalyst bed or, and a flash drum after the reactor separates the gas and liquid effluents. The gas phase, which is primarily hydrogen, is recompressed and recycled to the reactor inlet. When crude oil prices were low, there was little incentive to hydrogenate vacuum resid. Increased oil prices have made this process feasible, using an ebullated bed (slurry) hydrotreater. The reactor used to directly hydrogenate coal, shown in Figure 2.9, is essentially identical to the reactor used to hydrocrack vacuum resids. There are no industry-wide hydrogen purity standards for oil refining. Refiners prefer to use high purity hydrogen; but they can use lower purity hydrogen. The hydrogen produced by naphtha reformers typically contains substantial quantities of low molecular weight hydrocarbons. The recycled hydrogen used in hydrogenation typically contains low molecular weight hydrocarbons as well as H2S, NH3, and H2O, which are the products of sulfur, nitrogen, and oxygen removal from the oil feed. Impurities of all types lower the hydrogen partial pressure, which lowers the driving force for hydrogenation. Petroleum hydrogenation catalysts tolerate H2S and NH3, but relatively high levels of these components depress catalyst activity. Carbon monoxide and carbon dioxide in the fresh hydrogen feed are converted to methane and water via reactions R-7.14 and R-7.15.

COMBINED CYCLE PLANTS FOR POWER PRODUCTION, NGCC, AND IGCC Combined cycle plants burn gas or liquid fuels to generate electricity. When natural gas is the fuel, the process is known as natural gas combined cycle (NGCC). When syngas or syngas-derived gas is the fuel, the process is known as integrated gasification combined cycle (IGCC). Liquid fuels are generally much more expensive than gaseous fuels, so liquid fueled combined cycle plants are generally limited to locations where gaseous fuels are unavailable.

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Hydrogen Production and Integrated Gasiﬁcation Combined Cycle (IGCC)

Coal to liquids and coal to chemicals plants typically consume large quantities of electric power, especially for the air separation unit (ASU). These plants often generate power for internal consumption using a combined cycle plant. The fuel for these plants is either syngas, or a byproduct fuel gas such as the low molecular weight hydrocarbons produced by Fischer-Tropsch synthesis. A combined cycle plant is shown in Figure 7.5. Configuration details and operating conditions of these plants often vary. Much of the specific information shown here comes from Woods et al..1 Ambient air enters the compressor and is compressed to 1.6 to 1.9 MPa (16 to 19 compression factor). The compressor and the gas turbine usually have a common shaft. The compressor/turbine combination is known as a gen set, and can be broadly divided into two categories. Aero-derived gen sets are modified jet aircraft engines, and are built with light-weight, high-strength materials.

fuel

burner box electric generator stack

compressor

turbine hot flue gas

air

NH3 SCR cool flue gas HSRG

high pressure steam boiler feed water

high pressure steam turbine & generator intermediate pressure steam

condenser

intermediate pressure steam turbine & generator low pressure steam

exhaust steam

low pressure steam turbine & generator

Figure 7.5 Combined cycle plant. Steam from syngas processing is not shown.

Hydrogen Production and Integrated Gasiﬁcation Combined Cycle (IGCC)

Ground-based gen sets are specifically built for combined cycle plants, and are built with heavier, less expensive materials. Compressed air and fuel gas are fed to the burner box. The gases will leave the burner box at about 1,350 to 1,400 C. A low-NOx burner is used. Inert gases, generally steam in the case of NGCC, and N2 (from the ASU) in the case of IGCC, are often added to the burner box to add mass to the gas and to lower the flame temperature; which will greatly minimized the NOx emissions. Gas leaves the turbine at about 630 C and slightly above atmospheric pressure. Mechanical power produced by the turbine drives the compressor and an electric generator. In an NGCC plant, this generator produces about 65% of the electric power generated by the plant. Gas from the turbine enters the heat recovery steam generator (HSRG), which is a set of heat exchangers used to extract sensible heat from the flue gas to generate steam. The hot flue gas is initially used to generate high pressure steam at about 12.4 MPa. Flue gas, at about 260 to 340 C, is diverted to a selective catalytic reduction unit (SCR) to reduce NOx content (shown here as NO) via reaction with ammonia: 6NO þ 4NH3 /5N2 þ 6H2 O

R-7.16

The flue gas is then returned to the HSRG, where it is used to generate an intermediate pressure steam at about 2.9 MPa. Flue gas leaves the HSRG at about 105 C and is vented to the atmosphere though a stack. Gasification plants typically produce steam during syngas processing. This steam is then combined with the steam from the HSRG and fed to a set of steam turbines. The high pressure steam is fed to the high pressure steam turbine, and then returned to the HSRG for reheating. The intermediate pressure steam is fed to the intermediate pressure steam turbine. Exhaust from this turbine is fed to the low pressure steam turbine. The steam turbines are used to drive electric generators, which produce about 35% of the electric power from an NGCC power plant. Steam exhaust from the low pressure steam turbine is at sub-atmospheric pressure, typically about 7 kPa. This steam is sent to the condenser, and the condensate is returned to the HSRG as boiler feed water. Typically, cooling water from an evaporative cooling tower is used to cool the condenser. Evaporative loss from the cooling tower can be roughly 90% of the water consumed by an IGCC plant. Water consumption can be greatly reduced by replacing the evaporative cooling tower with air cooled condensers; but this will come with at higher capital cost.

NATURAL GAS COMBINED CYCLE (NGCC) NGCC plants were initially popular with power producers due to their low capital cost. For example, Woods et al.1 estimated that the capital cost of a nominal 500 MW power plant is $554/kW for an NGCC plant versus $1,549/kW for a subcritical pulverized coal

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Hydrogen Production and Integrated Gasiﬁcation Combined Cycle (IGCC)

plant. To meet rising power demand, producers built NGCC plants rather than the more expensive coal-fired and nuclear plants. The operating costs for NGCC plants are higher, however, because of the relatively high fuel cost. Consequently, NGCC power plants are run primarily as peaking units. When power demand is low, during the night, for example power is met primarily by coal-fired and nuclear power plants. Both of these plants which have relatively low fuel costs. They are run as base load units; which means that they are operated continuously at nearly maximum capacity. In the daytime, when power demand increases, NGCC plants are brought on line to meet the increased demand. At the time this book was written, a number of automobile manufacturers were developing battery-powered automobiles. These cars would be recharged during the night, and driven during the day. If battery-powered automobiles become popular, it could greatly change the daily fluctuation of power demand. NGCC power plants became popular due to their low CO2 emissions, compared to traditional pulverized coal power plants. Woods et al. estimated that the CO2 emissions from an NGCC power plant are 355 kg/MW-hr versus 807 kg/MW-hr for a subcritical pulverized coal plant burning Illinois No. 6 coal. Several US states set greenhouse gas emissions regulations for electric power based on typical NGCC emission rates. These regulations will become obsolete when US federal regulations require lower CO2 emission rates. Renewable energy power plants, such as solar and wind, produce no CO2 emissions, but the power output is variable, depending on whether the sun shines or the wind blows. Coupling a renewable energy plant with an NGCC plant would yield a more consistent, deployable power source. The NGCC plant would operate when insufficient power is available from the renewable energy plant. This combination would have lower CO2 emissions than if power were solely produced by the NGCC plant.

INTEGRATED GASIFICATION COMBINED CYCLE (IGCC) Integrated gasification combined cycle (IGCC) with carbon capture and sequestration (CCS) is an attractive means of generating electric power from coal with minimal greenhouse gas emissions. The basic approach is to convert most of the syngas to H2 and CO2 using the water gas shift reaction, and then separating the H2 and CO2 using an AGR. The CO2 is sequestered. The H2 is burned in a gas-fired turbine to produce electricity. IGCC offers substantial advantages over pulverized coal combustion when CCS is required. Removal of CO2 from pressurized syngas is much easier and less expensive than removing CO2 from a near atmospheric pressure flue gas produced by a pulverized coal plant. Woods et al.1 estimated that the capital costs for IGCC with CCS were $2,390/kW to $2,668/kW versus $2,895 for a subcritical pulverized coal combustion plant with CCS. This was based on a 90% carbon capture level, and the resulting IGCC

Hydrogen Production and Integrated Gasiﬁcation Combined Cycle (IGCC)

CO2 emissions were 68 to 86 kg/MW-hr, only 19 to 24% of the emissions of an NGCC plant without CCS. Even with this capital cost advantage, CCS is an expensive technology. The most attractive IGCC with CCS plant had a levelized cost of electricity that was 60% higher than the cost of electricity from a traditional pulverized coal combustion plant. IGCC without carbon capture and sequestration (CCS), has little to offer compared to traditional, subcritical pulverized coal combustion. Woods et al. estimated higher capital costs, leading to a higher cost for electric power production. Due to higher efficiency, CO2 emissions for IGCC plants are a little lower than for pulverized coal plants, but still nearly double the emissions of an NGCC plant. Although the technology for IGCC is similar to NGCC, the economics differ greatly. Syngas production and processing greatly increases the capital cost; and coal costs much less than natural gas. Consequently, IGCC plants are more appropriate as base load units. At the time this book was written, there were no fixed standards for H2 purity for IGCC. Greenhouse gas regulations were in their infancy, and fixed emissions targets were not set. The FutureGen plant21 will initially target a 60% carbon capture rate, which will increase to 90% by its third year of production. The similar ZeroGen22 plant in Queensland, Australia, plans a 65% initial carbon capture rate, which will increase to a 90% carbon capture rate. Both the FutureGen plant and the design studies by Woods et al. are based on a single-stage sour gas shift reactor followed by a Selexol AGR plant. Several US states set greenhouse gas regulations for purchased electric power. Current carbon dioxide limits in these regulations approximately correspond to the CO2 emissions for NGCC. If greenhouse gas regulations follow the trend previously observed for other air emissions; the standards will become increasingly stringent over the next few decades. Following this pattern, power producers will initially meet CO2 targets by burning more natural gas and less coal. When NGCC without CCS can no longer meet the CO2 emission targets, IGCC with CCS may be an attractive option. In the study by Woods et al., syngas left the sour gas shift reactor at 270 C, and the gas fed to the turbine contained 91.0% H2, 1.9% CO, 4.5% CO2, and 0.08% CH4. Much of the CO2 in the turbine feed gas came from a CO2 purification process off-gas stream. Carbon capture rates greater than 90% are possible by adding an LT shift reactor, and by routing the off-gas stream back to the Selexol unit. If the hydrogen produced from IGCC met the purity standards currently used for ammonia, carbon capture rates greater than 99.5% are possible, although the cost per ton of carbon dioxide captured would be considerably higher. Fluidized bed and moving bed gasifiers would produce syngas with greater hydrocarbon content; and these hydrocarbons would burn to produce CO2. To meet the same CO2 emission targets, fluidized bed and moving bed gasifiers would have to push the water gas shift reaction further to the right.

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Hydrogen Production and Integrated Gasiﬁcation Combined Cycle (IGCC)

If the cost of natural gas makes peak load generation with NGCC prohibitive, or if NGCC is no longer able to meet greenhouse gas regulations, another source of peak power will be needed. Before NGCC became popular, the usual approach was to install sufficient coal-fired generation capacity to meet peak load demands. This is a workable, but expensive approach, especially if IGCC with CCS replaces pulverized coal combustion. One idea to lower this cost would be to use large batteries to store electric power during low demand periods, and release that stored energy during peak load periods. Most recent battery development seeks high energy/weight ratios for portable electronics and vehicles. For utility use, the battery design needs to emphasize low cost and durability. Another concept would be to produce a storable turbine fuel using a gasification plant, and then burn that stored fuel during peak power demand. For example, H2 could be stored using one of the techniques described in Chapter 8.

COMBINING IGCC AND OXY-COMBUSTION In Chapter 2, the oxy-combustion concept was described. Instead of air, a pulverized combustion plant is fed a blend of O2 from an air separation unit (ASU) and recycled CO2. The flue gas can be compressed and sequestered without separation. A similar concept can be applied to IGCC, as shown in Figure 7.6. Hot syngas, with particulates removed but otherwise unprocessed, is fed to the gas turbine along with O2 from an ASU. Recycled CO2 flue gas replaces N2 from air to provide gas mass for the turbine and to achieve the proper flame temperature. Hot flue gas leaving the turbine is sent to an HSRG to recover sensible heat. The cool flue gas leaving the HSRG is fed to a fan. Part of the output of the fan is recycled to the gen set, and the remainder is compressed and sequestered.

steam turbines/ generators/ condenser O2 hot syngas

compressor/ gas turbine/ generator

HSRG

fan recycle CO2

Figure 7.6 A combination of oxy-combustion and IGCC with CCS.

CO2 to compression & sequestration

Hydrogen Production and Integrated Gasiﬁcation Combined Cycle (IGCC)

An advantage to this approach is that the cost and energy penalty of syngas processing is eliminated. The syngas fed to the burn box is hot, instead of the cool hydrogen used in the usual approach. A disadvantage is that a larger ASU would be required to replace O2 in the air fed to the compressor.

METHANOL, SNG, AND FISCHER-TROPSCH SYNTHESIS For the previously considered applications, hydrogen is the intermediate gas product; and the other components in the gas stream are regarded as impurities. Methanol, SNG, and Fischer-Tropsch (FT) synthesis differ in that a blend of H2 and CO is desired. The water gas shift reactor is operated to achieve the desired H2/CO ratio, rather than maximize H2 conversion. In the process and economic study by Tomlinson et al.,12 the steam-iron process was used to convert a portion of the desulfurized syngas to nearly pure H2. This stream was then blended with the remainder of the syngas to achieve the desired H2/CO ratio for Fischer-Tropsch synthesis.

REFERENCES 1. Woods MC, Capicotto PJ, Haslbeck JL, et al. Cost and performance baseline for fossil energy plants, volume 1: Bituminous coal and natural gas to electricity. DOE/NETL-2007/1281, <www.netl.doe. gov/energy-analyses/pubs/Bituminous%20Baseline_Final%20Report.pdf; 2007.> 2. Southern company services, Power Systems Development facility Topical Report, Test campaign TC-16. U.S.D.O.E. contract_DE-FC21-90MC25140, <www.netl.doe.gov/technologies/coalpower/gasification/ projects/adu-gas/25140/TC-16-TopicalReport2004.pdf>. 3. Bartone Jr LM, White J. Industrial size gasification for syngas, substitute natural gas, and power production, DOE/NETL-401/040607, <www.netl.doe.gov/energy-analyses/pubs/Final%20Report %20DOE_NETL-401_1.zip; 2007>. 4. Lloyd L, Ridler DE, Twigg MV. The water-gas shift reaction. In: Twigg MV, ed. Catalyst Handbook. 2nd ed. Wolfe Publishing; 1989. p. 283-339. 5. Newsome DS. The water-gas shift reaction. Catal Rev-Sci Eng. 1980;21:275-318. 6. Podolski WF, Kim YG. Modeling the water-gas shift reaction. Ind Eng Chem Process Des Develop. 1974;13:415-421. 7. Bohlbro H. The kinetics of the water gas conversion: III. The influence of H2S on the Rate Equation. Acta Chem Scand. 1963;17:1001-1011. 8. Pearce BB, Twigg MV, Woodward C. Methanation. In: Twigg MV, ed. Catalyst Handbook. 2nd ed. Wolfe Publishing; 1989. p. 340-383. 9. Haldor Topsoe, Inc. Sulfur resistant/sour water-gas shift catalysts. product brochure; 2009. 10. Hurst S. Production of hydrogen by the steam-iron method. J Am Oil Chem Soc. 1939;16:29. 11. Wenguo X, Yingying C. Hydrogen and electricity from coal with carbon dioxide separation using chemical looping reactors. Energy & Fuels. 2007;21:2272-2277. 12. Tomlinson G, Gray D, White C. Chemical looping process in a coal-to-liquids configuration, independent assessment of chemical-looping in the context of a Fischer-Tropsch Plant. DOE/NETL-2008/ 1307 (2007). <www.netl.doe.gov/energy-analyses/pubs/DOE Report on OSU Looping final.pdf.> 13. Plunkett J, Gray D, White C, et al. Performance and cost comparisons of alternate IGCC based CO2 capture technologies. Gasification Technologies; 2009. Annual Conference. 14. Gupta P, Velazquez-Vargas LG, Fan L-S. Syngas redox (SGR) Process to produce hydrogen from coal derived syngas. Energy & Fuels. 2007;21:2900-2908.

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15. Li F, Kim HR, Sridhar D, et al. Syngas chemical looping gasification process: Oxygen carrier particle selection and performance. Energy & Fuels. 2009;23:4182-4189. 16. Gasior SJ, Forney AJ, Field JH, et al. Production of synthesis gas and hydrogen by the steam-iron process: Pilot Plant study of fluidized and free-falling beds. U.S. Bureau of Mines Report of Investigations. 1961;5911. 17. Wakker JP, Gerritsen AW, Moulijn JA. High temperature H2S and COS removal with MnO and FeO on g-Al2O3 acceptors. Ind Eng Chem Res. 1993;32:139-149. 18. Jennings JR, Ward SA. Ammonia synthesis. In: Twigg MV, ed. Catalyst Handbook. 2nd ed. Wolfe Publishing; 1989. p. 384-440. 19. Thomas G. Overview of storage development DOE hydrogen program. Hydrogen Program Review; 2000. 20. Arienti S. Gasification to meet refinery hydrogen, electricity, and steam demands: Availability vs. costs. Gasification Technologies; 2009. Annual Conference. 21. Brown JD. Rebirth of FutureGen at Mattoon. Gasification Technologies; 2009. Annual Conference. 22. Grieg C. ZeroGen project update: Commercial scale IGCCþCCS demonstration, managing risks and uncertainties. Gasification Technologies; 2009. Annual Conference.

CHAPTER

8

Hydrogen Adsorption and Storage Xin Hu, Maohong Fan, Brian Francis Towler, Maciej Radosz, David A. Bell, Ovid Augustus Plumb Contents Introduction Physisorption of Hydrogen Hydrogen Adsorption in Porous Materials e Excess Adsorption Versus Total Uptake Theoretical Hydrogen Storage Investigations of Nanostructured Carbon Materials Experimental Results of Carbon Nanostructured Materials Initial Promising Results and Discrepancies Comparative Hydrogen Storage Studies in Different Carbon Nanostructures Hydrogen Storage in Carbon Nanotubes Hydrogen Storage in Graphite/Carbon Nanoﬁbers Hydrogen storage in activated carbon Hydrogen Storage in Other Carbon Nanostructures Hydrogen Storage in Carbon Nanostructures by Spillover

Hydrogen Storage of Metal-Organic Frameworks Hydrogen Storage of Zeolites Hydrogen Storage of Clathrates Chemisorption of Hydrogen Hydrogen Storage of Metal Hydrides Hydrogen Storage of Complex Hydrides Alanates Borohydrides

Hydrogen Storage in Pure And Reduced Microporous and Mesoporous Ti Oxides References

157 158 159 160 166 166 168 173 181 183 184 187

188 204 207 210 210 218 218 226

230 235

INTRODUCTION As fossil fuels become more limited, the demand for an efficient and clean fuel alternative has increased, and is expected to become even more pronounced in the future. Hydrogen represents an ideal alternative as a fuel as it exhibits the highest heating value per mass of all the chemical fuels. Furthermore, it is the most abundant element in the universe, 1 it can be readily synthesized, and when used in a fuel cell with oxygen it produces only water as a byproduct, so it is environmentally friendly. However, the main concern to date is the safe and efficient transport of this extremely flammable gas. Two different storage technologies are conventionally used, high-pressure gas cylinders and liquid hydrogen in Coal Gasification and Its Applications. ISBN B978-0-8155-2049-8.10008-7, doi:10.1016/B978-0-8155-2049-8.10008-7

Ó 2011 Elsevier Inc. All rights reserved.

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cryogenic vessels.2 But the former suffers from low storage density and safety concerns due to the high pressures required; while the large amount of energy required for liquefaction and continuous boil-off of liquid hydrogen are significant drawbacks for the latter. To overcome these problems, solids which absorb hydrogen reversibly with high gravimetric and volumetric density, and operate under moderate temperature and pressure may be the most desirable solution. In spite of the recent surge in interest, hydrogen storage has been a materials science challenge for decades. The goal is to pack hydrogen as close as possible to achieve the highest volumetric and gravimetric density. The second important criterion for a hydrogen storage system is the reversibility of uptake and release. The US Department of Energy (DOE) has set 9.0 wt.% and 81 kg H2/m3 as the 2015 system targets. The main challenges in the field of hydrogen storage are to devise new materials that (1) are inexpensive to manufacture in large quantities in a pure form, (2) have high gravimetric and volumetric hydrogen storage density, and (3) exhibit fast sorption kinetics at near-ambient temperature and have a high tolerance for recycling. One of the most promising classes of materials for hydrogen storage are nanostructured composites, as they have dramatically different chemical, physical, thermodynamic and transport properties compared to their bulk counterparts. Due to the wide range of compositions, the ability to tailor pore and grain sizes, and the capacity to intimately weave two or more phases together at the nanometer level, nanophase composite materials may open the window to greater hydrogen storage capacities and lower kinetic adsorption barriers. This is compared to the coarse grained materials that are currently available. Generally, hydrogen storage materials can be divided into two categories: those that bind molecular hydrogen to the surface via weak dispersive interaction in a process known as physisorption, and those that trap atomic hydrogen through a strong chemisorption process (e.g. metal hydrides, complex hydrides). Both types of storage methods will be the subject of this chapter.

PHYSISORPTION OF HYDROGEN Physisorption is a principle where the weak intermolecular forces, which do not cause any significant change in the electronic orbital patterns, are involved. These forces are also known as van der Waals forces. A combination of long-range attractive, dispersive or van der Waals, interactions and short-range repulsive interactions between the adsorbent and the adsorbate molecules are responsible for this phenomenon. This results in a minimum in the potential energy curve of the gas at approximately one molecular radius from the solid surface. The attractive interaction originates from long-range forces produced by fluctuations in the charge distribution of the gas molecules and of the atoms on the surface. This gives rise to a net attraction between temporary fluctuating and induced dipoles. However, at small distances, the overlap between the electron cloud of the gas molecule and of the substrate is significant and the repulsion increases rapidly.

Hydrogen Adsorption and Storage

Once a monolayer of adsorbate molecules or atoms has formed, the gaseous species interacts with the liquid or solid adsorbate. Therefore, the binding energy of the second layer of adsorbates is similar to the latent heat of sublimation or vaporization of the adsorbate. Consequently, adsorption at a temperature at or above the boiling point of the adsorbate at a given pressure leads to the adsorption of a single monolayer. Compared with hydrogen storage in metal hydrides and complex hydrides, physical adsorption has the advantage of being completely reversible and of exhibiting very fast kinetics. In addition, since a very small amount of energy (< 10 kJ mole1) is involved both in the adsorption and the release of H2, no extra heat management systems are required for onboard applications. On the other hand, due to of the low adsorption enthalpy involved in physisorption, low temperature typically is necessary to reach high storage capacities. Carbon nanostructured materials3e10 have received the most attention in this area due to their low density, high surface area, extensive pore structure, wide variety of structure forms, good chemical stability and amenability to a wide range of preparation, carbonization, and activation conditions. Newly found carbon nanostructures include activated carbon, carbon nanotubes, and carbon nanofiber. An overview of theoretical and experimental aspects of hydrogen storage in carbon materials follows.

Hydrogen Adsorption in Porous Materials e Excess Adsorption Versus Total Uptake There are two ways to quantify the extent of hydrogen adsorption, they are the excess adsorption and the total uptake. In a recent review,11 Long and his co-workers defined the excess adsorption as the amount of H2 taken up beyond what would be contained within free volume that is equivalent to the total pore volume of the sorbent under identical conditions; free volume is the volume unoccupied by the sorbent atoms and molecules. Thus, this quantity approximates the amount of H2 adsorbed on the surface alone. Since the efficiency of packing gas molecules within the confines of the pores is lower than that achieved in free volume itself, the excess adsorption will increase in pressure until it reaches a maximum (say at 20e40 bar), and then it will decrease. Despite the decrease, excess adsorption data at pressures above the maximum are of value for assessing the compressibility of H2 within the material and evaluating the total uptake. The total uptake, sometimes referred to as the absolute uptake, corresponds to the amount of hydrogen contained within the boundaries formed by the adsorbent. This quantity includes both surface-adsorbed H2 and the H2 gas compressed within the pores. To calculate the total uptake from the excess adsorption, it is necessary to know the precise density of the material skeleton, or the free volume of the adsorbent. This is typically measured using helium gas. The total uptake can be used to estimate the volumetric storage density, which is often used to select hydrogen storage materials. It is important to note, however, that the total uptake does not account for the efficiency of packing the sample in a container, which determines the bulk density and hence the size

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of the storage system.11 At low pressure, say at 1 bar, the excess adsorption and the total uptake are nearly identical. However, at higher pressure, these two values can differ significantly. Unless stated otherwise, all hydrogen storage data discussed in this chapter will be the excess adsorption.

Theoretical Hydrogen Storage Investigations of Nanostructured Carbon Materials Many theoretical studies have been performed to predict the hydrogen sorption in carbon nanostructured materials. This is especially true for CNTs (carbon nanotubes), utilizing GCMC (grand canonical Monte Carlo) simulations.12e20 Rzepka et al.12 investigated the physisorption of hydrogen molecules on the surface atoms of carbon slit pores and CNTs using GCMC calculations. The brief description of the simulation is described as follows. In a simulation loop, the motion of hydrogen molecules in a given pore volume for a fixed temperature T and a chemical potential m is calculated. The carbon pore is built up by M carbon atoms (~2000) located at the surface of any desired pore geometry. Two geometries shown in Figure 8.1, slitpores and nanotubes, are used in this calculation. As indicated, the carbon atoms are arranged in a hexagonal pattern inside the walls of the slitpore. To allow a continuous variation of the tube diameter d, a rectangular pattern at the circumference of the nanotubes was chosen. A test of the influence of the arrangement of atoms inside the pore walls for a carbon slitpore shows no significant difference between a hexagonal and rectangular pattern. ˚ 2 for both geometries. The The particle density in the pore walls is Ac ¼ 0.382 A pairwise interaction energy between two particles separated by a distance s is calculated by the Lennard-Jones potential s s UðsÞ ¼ 43½ð Þ12 ð Þ6 s s

Eqn. 8.1

Figure 8.1 Pore geometries studied in this work: slitpore with pore size d and cylindrical pore (nanotube) with diameter d. The arrangement of individual carbon atoms inside the pore walls is as indicated.12

Hydrogen Adsorption and Storage

˚ and 31/kB ¼ 33.3 K (hydrogen-hydrogen) and The potential parameters, s1 ¼ 2.97 A ˚ s2 ¼ 3.19 A, 32/kB ¼ 30.5 K (hydrogen-carbon) are explored in this study. The total potential energy of a particular hydrogen molecule is finally given by summing all interaction energies (Eqn. 8.1) between neighboring hydrogen molecules and pore wall carbon atoms. During the simulation the particle number N within the pore fluctuates due to particle displacement, creation, and destruction steps, which are executed with equal frequency. In a creation step, the position of the new particle is chosen randomly within pore volume V and its potential energy U is calculated. Finally this creation step is accepted with the probability min(1; P) with U m V P ¼ exp ðcreationÞ Eqn. 8.2 kB T N þ 1 Analogous destruction and displacement steps are accepted with U m N ðdestructionÞ P ¼ exp þ kB T V U1 U2 P ¼ exp ðdisplacementÞ kB T

Eqn. 8.3

Eqn. 8.4

To obtain thermal equilibrium, 106 iterations were performed. Then an additional 105 iterations were performed to obtain a simulation result for the mean particle number hNi in the pore. This is in the range between 0 and 500 for the parameters used in this calculation. Furthermore, values for the following commonly used parameters were calculated: mH2 Eqn. 8.5 Volumetric storage capacity ¼ V Gravimetric storage capacity ¼ Excess adsorption ¼

mH2 mH2 þ mC

mH2 moH2 ðmH2 moH2 Þ

Eqn. 8.6

Eqn. 8.7

With M $0:5$12:011g=mol L

Eqn. 8.8

mH2 ¼

hN i $2:016g=mol L

Eqn. 8.9

moH2 ¼

N0 $2:016g=mol L

Eqn. 8.10

mC ¼

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The factor 0.5 in the carbon mass considers a periodic arrangement of the single pores in the whole sample in which each pore wall joins two micropores on both wall sides. N0 gives the number of hydrogen molecules corresponding to the density of compressed gas inside the accessible pore volume. The relation between the chemical potential m and the pressure p at a given temperature is calculated by a similar GCMC simulation of bulk hydrogen gas. The simulation results showed that at room temperature and 10 MPa, a gravimetric storage capacity of 1.3 wt.% could be achieved for the optimum pore geometry. A slit pore consists of two graphitic layers separated by a distance that corresponds to approximately twice the diameter of a hydrogen molecule (~ 7 A˚), compared with 0.6 wt.% for the nanotube geometry. These calculated hydrogen storage capacities are consistent with the experimental results. In addition, at nearly all temperatures and pressures the storage capacity of slit pores is better than the capacity of carbon nanotubes. Later, Williams and Eklund13 performed GCMC simulation of H2 physisorption in finitediameter ropes of carbon SWNTs, or single-walled carbon nanotubes. The strong dependence of the gravimetric adsorption on the diameter of a SWNT rope was found to be correlated with computed values of the specific surface area. Furthermore, the simulations showed adsorption energies for interstitial channel of 11 kJ mol1, followed by the grooves (9 kJ mol1) and endohedral pore (6 kJ mol1) inside the tube, and the sites on the outer surface of the bundles with only 5 kJ mol1 (Figure 8.2). Thus, they suggested that delamination of nanotube ropes should increase the gravimetric storage capacity. A maximum hydrogen storage capacity of 9.6 wt.% was calculated for an isolated (10, 10) nanotube at 77 K and 10 MP, compared with 7.0 and 5.5 wt.% for the 3- and 5-tube ropes under the same conditions, respectively. Wang and Johnson14 calculated the hydrogen adsorption isotherms in SWNT arrays, isolated SWNTs and idealized carbon slit pores, respectively. They compared the simulations with experimental data for the activated carbon AX-21. The idealized slit pores gave significantly better performance for hydrogen storage than SWNT arrays, whereas the gravimetric density of hydrogen in isolated tubes at 77 K and 50 atm was well above that surface

interstitial channels

groove endohedral pore

Figure 8. 2 . Different storage sites on a rope of SWNTs.

13

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for the activated carbon AX-21 and the idealized slit pore. This indicated that the packing geometry of the SWNTs played an important role in hydrogen storage. The results showed that if tubes in the array could be separated from one another, the sorbent would have a greater effective surface area and volume available for adsorption. Furthermore, they15 optimized the packing geometry of SWNTs to enhance their hydrogen uptake abilities. It was found that the (9, 9) tube triangle arrays with an ˚ had the maximum volumetric density of 10.1 and 60 kg/m3 intertube distance of 6 and 9 A at 298 K, 10 MPa and 77 K, 10 MPa. In another study, Wang and Johnson16 calculated the hydrogen adsorption for GNFs, or graphitic nanofibers. The maximum hydrogen uptake of 1.6 wt.% was obtained for a GNF with a pore width of 9 A˚ at 298 K and 50 atm. Darkrim and Levesque17 computed the hydrogen adsorption in two arrangements, both square and hexagonal, of open SWNTs by Monte Carlo simulations. The simulation predicted maximal hydrogen adsorption of 11.24 wt.% in gravimetric and 60 kg/m3 in volumetric capacity for the SWNT with a diameter of 2.2 nm, intertube spacing of 1.1 nm and a square arranged lattice at 77 K and 10 MPa. These results are in agreement with the experimental data for SWNT by Ye and his co-workers.21 Yin et al.18 performed GCMC simulations on hydrogen storage in a triangular array of SWNTs and in slit pores at 298 K and 77 K, respectively. The maximum hydrogen adsorption capacity at 298 K was 4.7 wt.% for the SWNTs with a diameter of 6.0 nm and intertube spacing of 1.0 nm at 100 bar, whereas at 77 K and 70 bar, the maximum hydrogen uptake of 33 wt.% was evaluated for the SWNTs with a diameter of 6.0 nm and intertube spacing of 3.0 nm. They claimed that the adsorption occurred entirely in the interstitial space between the nanotubes. However, such a high hydrogen storage capacity has not been confirmed experimentally at this time. For slit pores, the hydrogen uptake at both temperatures was lower than those of SWNTs. More recently, Guay et al.19 investigated the hydrogen storage behavior of SWNTs, DWNTs (doublewalled carbon nanotubes) and GNFs using the GCMC method. The amount of hydrogen uptake was strongly influenced by the structure porosity. A maximum capacity of 1.4 wt.% was achieved for the nanostructured carbons with optimum pore diameter of around 0.7 nm at 293 K and 10 MPa. Furthermore, they also predicted that pure carbon nanostructures could not reach the DoE target at room temperature based on their simulation results. Volpe and Cleri20 studied the hydrogen sorption in chemically modified matrices of SWNTs by using a combination of tight-binding molecular dynamics and GCMC. A maximum hydrogen uptake capacity of 3.4 wt.% was achieved ˚ arranged in a square lattice with for the SWNTs with an optimized diameter of 11.7A ˚ a wall-to-wall intertube distance of 7 A at 293 K and 10 MPa. They proposed that surface modification by oxidation or electron, or ion-beam irradiation could improve the hydrogen storage capacity, exploiting the combined effect of surface curvature and maximum specific surface exposed to the gas. In order to study the formation of carbon-hydrogen bonds, DFT (density functional theory) and ab initio calculations have

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been performed by several groups recently. Lee et al.22 have reported results of calculation for hydrogen storage behavior in SWNTs by DFT calculations at 0 Kelvin, and found two chemisorption sites at top sites of the exterior and the interior of the tube wall. They further found that a form of H2 molecule could exist in an empty space inside nanotubes, and that the storage capacity of hydrogen in this empty space increased linearly with tube diameter. The maximum storage capacity was limited by the repulsive energies between H2 molecules inside nanotubes, and those between H2 molecules and the tube wall. Their calculation predicted that the hydrogen storage capacity in (10, 10) nanotubes can exceed 14 wt.% and 160 kg/m3. Later that year, they23 published a comprehensive study on hydrogen storage in SWNTs and MWNTs (multi-walled carbon nanotubes) by performing density-functional-based tight-binding calculations. It was found that the storage capacity for hydrogen increased linearly with tube diameters in SWNTs. A maximum uptake of 14 wt.% can be predicted, whereas hydrogen storage capacity was independent of diameter in MWNTs and a maximum uptake of 7.7 wt.% was estimated. They confirmed, using a Raman spectra, that hydrogen existed as a form of H2 molecule in an empty space inside CNTs. In a further work, they investigated the hydrogen-adsorption mechanism in the (5, 5) SWNTs with DTF calculations at zero temperature.24 The hydrogen atoms first adsorbed on the tube wall, and then were stored in the capillary as a form of H2 molecule at higher coverage. Instead of the capillary effect through the ends of the nanotubes, the H2 can be stored in the capillary through the tube wall by flip-in and kick-in mechanism, while preserving the wall stability of a nanotube. The authors claimed that this calculation may describe an electrochemical storage process for hydrogen, which is applicable for the secondary hydrogen battery. Moreover, Ma et al.25 studied the hydrogen-storage behaviour of (5, 5) SWNTs using molecular dynamics simulations and ab initio electronic calculations. Hydrogen atoms with kinetic energy of 16e25 eV were found to penetrate and be trapped into the tube in high-density liquid hydrogen form (Figure 8.3). The hydrogen storage capacity for a (5, 5) SWNT can be 5 wt.% and 132.4 kg/m3 in gravimetric and volumetric density, respectively. Zhang et al.26 studied the adsorption of hydrogen on SWNTs with different diameters by a combination of the classical potential and DFT methods. A maximum value of around 1 wt.% was obtained at 20 MPa for the isolated SWNT sample at room temperature. At 77 K and 4 MPa, maximum gravimetric capacities of 7.1 wt.% and 9.5 wt.% could be achieved for bundled and isolated SWNTs with diameter of 2.719 nm respectively. The hydrogen uptake results for bundled SWNTs are in good agreement with previous experimental results.21 In another work, Bauschlicher et al.27 calculated the binding energy of H to a (10, 0) CNT at different H coverage using the AM1 and ONIOM approaches. It was found that the 50% H coverage on the carbon nanotube was the most favorable. This corresponds to about 4 wt.% hydrogen storage capacity. Cheng et al.28 performed analysis of quantum-mechanical

Hydrogen Adsorption and Storage

Figure 8.3 Hydrogen stored in a (5, 5) SWNT capsule via H atom implantation at 20 eV. (a) The side view of the structure, showing H2 molecules formed from the injected H atoms; (b) the top view, showing H2 molecules condensed to a single shell of tube-shaped liquid.25

molecular dynamics simulations of hydrogen binding energies in (9, 9) SWNTs, of which 7.51 and 6.75 kcal/mol were obtained at 300 K for the inside tube and outside tube configuration, respectively. They argued that partial electron-transfer interaction between hydrogen and instantaneously distorted carbon atoms in the SWNT wall was the primary reason for the high hydrogeneSWNT adsorption energy (Figure 8.4). Han et al.29 investigated the interaction of H2 on the exterior surface of single-walled and bundle CNTs by using molecular dynamics simulations. An interesting phenomenon of hydrogen gas molecules transformed into a liquid phase on the surface of a CNT bundle was observed at 80 K and 10 MPa. This is due to the long-range electrostatic interaction of polarized charges on the deformed CNT bundles with the H2 molecule. This does not occur on a single CNT, but it is a result of its symmetrical deformation under the same pressure. This liquefaction can account for the higher hydrogen storage capacity on the CNT bundle rather than on the CNT. The authors also claimed that the H2 gas might liquefy at temperature higher than 80 K on a more strongly polarized CNT bundle. Lee

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Figure 8.4 A representative time step during the MD simulation of endohedral H2 adsorption at 300 K. The radial deformation of the SWNT wall and orientation of the H2 molecules are illustrative of the dynamics observed throughout the simulation.28

et al.30 reported the hydrogen storage capacity and desorption mechanisms of Ni-dispersed CNTs by DFT methods and experimental methods. They determined that each Ni dispersed on the surface of CNTs could store up to five hydrogen molecules with an enthalpy change of 0.26 eV/H2 in hydrogen adsorption. This is consistent with the 0.32 eV/H2 from experiment (Figure 8.5). Furthermore, the maximum H2 uptake of 10 wt.% was predicted for a high Ni-coverage, and Ni-dispersed SWNT at room temperature. From the above theoretical results, it is clear that both positive13,15,17,18,23,24,30 and negative12,19 results had been obtained. Some simulation results are supported by experimental data,17,26 but others not,18,23,24 possibly due to the improper selection of calculation functions. However, these theoretical studies have provided useful information for the design of optimal carbon nanostructures for hydrogen storage.

Experimental Results of Carbon Nanostructured Materials Initial Promising Results and Discrepancies In 1997, Dillon et al. first reported hydrogen storage properties for SWNTs. They measured the hydrogen uptake of a small quantity of soot containing 0.1e0.2 wt.% SWNTs at room temperature using a TPD, temperature programmed desorption, spectroscope.31 Hydrogen adsorption capacity for pure SWNTs was estimated in the range of 5e10 wt.%. A high heat of adsorption of 19.6 kJ/mol was also obtained. The authors claimed that hydrogen can condense to high density inside narrow SWNTs, and predicted that SWNTs with a diameter of 1.63 and 2 nm would come close to the target hydrogen uptake density of 6.5 wt.% set by the DoE. Later, they32 reported hydrogen storage

Hydrogen Adsorption and Storage

(a)

(b)

Figure 8.5 (a) H2 molecules adsorbed around Ni dispersed on the surface of a carbon nanotube. (b) View for molecular orbitals for 5 H2 molecules adsorbed on the CNT.30

capacity of about 7 wt.% after purifying the samples and opening the SWNTs. However, Hirscher et al.33 showed that the Ti alloy particles introduced in the sample during the ultrasonic treatment were responsible for most of the hydrogen storage capacity of SWNTs. In further studies, they34,35 reported a low hydrogen storage capacity of 1.0 wt.% for purified SWNTs, 0.5 wt.% for GNFs and 0.3 wt.% for graphite. They noted that the positive results reported for hydrogen storage in CNTs had not been repeated or confirmed independently at other laboratories. More recent reports by Heben’s group36,37 showed significantly lower hydrogen storage capacities of 2e3 wt.% for SWNTs. In 1999, Ye et al.21 reported a hydrogen storage capacity of 8 wt.% for crystalline ropes of SWNTs at a cryogenic temperature of 80 K and pressure higher than 12 MPa. They suggested that the tube bundle structure expanded under high pressure to enable higher adsorption on newly uncovered surface. Liu et al.38 reported that a hydrogen storage capacity of 4.2 wt.% could be reached at room temperature and 10 MPa for SWNTs synthesized by the hydrogen arc discharge method and with a larger mean diameter of 1.85 nm. However, about 20% of the absorbed hydrogen remained in the sample after desorption at room temperature. Even more promising, Chambers et al.39 reported extremely high hydrogen storage capacities of up to 67 wt.% for a herringbone-type GNF at room temperature and 110 atm. In a further study40 they reported results on the interaction of hydrogen with GNFs, and proposed that the high hydrogen uptake was due to the special structural conformation of GNF. This can produce a material composed entirely of nanopores that accommodate the H2 molecules. They also pointed out that the pretreatment of samples before hydrogen storage is very important. These extremely high storage capacities, however, have not been independently confirmed. For example, Ahn et al.41 found at a pressure of 8 MPa at 77 K and

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18 MPa at 300 K less than 0.01 wt.% hydrogen storage in GNFs applying the volumetric method. In another work, Tibbetts et al.42 examined different carbon materials (e.g. nanotubes, carbon fibers, high surface area activated carbons) at high pressure and they found the maximum hydrogen uptake for all the samples was less than 0.1 wt.% at room temperature and 3.5 MPa. They claimed that hydrogen storage capacity higher that 1 wt.% for carbon materials at room temperature was due to experimental error. Ritschel and co-workers43 also investigated hydrogen adsorption on different carbon nanostructure materials at room temperature and pressure up to 45 bar. The measurements showed limited hydrogen storage capacity of less than 0.7 wt.% for all the samples. Moreover, Lueking et al.44 studied the microstructure and hydrogen storage properties of various CNFs synthesized from different carbon sources and catalysts, and then subjected to various pre-treatments. The CNF possessing a herringbone structure and a high degree of defects exhibited the best performance for hydrogen storage, and a maximum hydrogen desorption capacity of 3.8 wt.% was achieved at 69 bar and room temperature. However, this value is much lower than the result previously reported.39 In 1999, Fan et al.45 investigated the hydrogen absorption of vapor-grown CNFs with a diameter of 5 nm to 300 nm. The maximum hydrogen uptake of 12.38 wt.% was obtained at room temperature and 110 atm. The nanofibers had to be boiled in hydrochloric acid before the hydrogen storage test. However, in a further publication, the same group46 reduced the storage capacity of the CNFs by a factor of two. In the same year, Chen et al.47 reported that high hydrogen uptakes of 20 wt.% and 14 wt.% could be achieved in Li-doped and K-doped MWNTs, respectively, at a pressure of 1 atm. However, their results have been questioned by Yang.48 Yang repeated their experiments and pointed out that it was the moisture in the hydrogen gas that drastically increased the weight gain by reactions with, or adsorption on, the alkali species on carbon, while the contribution of hydrogen storage might be limited. This conclusion has been supported by Skakalova and co-workers.49 In addition, Pinkerton et al.50 investigated the hydrogen storage properties of Li-doped CNTs and Li- and K- intercalated graphite. They found the K-intercalated graphite showed hydrogen adsorption of 1.3 wt.% only, while no hydrogen adsorption can be observed for the Li-containing carbon samples. This discrepancy in the hydrogen storage abilities of carbon nanostructures is considered to be due to difficulties in making accurate measurements, impurities in the sample, and poor understanding of the hydrogen sorption mechanism. Since then, scientists have made strong efforts to identify factors such as measurement methodology, material structure characterization, and synthesis techniques that influence the hydrogen storage capacity of carbon nanostructures. Subsequently, both positive and negative results have been reported. Comparative Hydrogen Storage Studies in Different Carbon Nanostructures In 2001, Nijkamp et al.51 reported the hydrogen storage capacities for a wide variety of carbonaceous sorbents, microporous zeolites, and non-porous materials at 77 K and

Hydrogen Adsorption and Storage

pressure up to 1 bar. Reversible physisorption took place exclusively in all samples. An approximately linear relationship between the hydrogen uptake capacity and the surface area of samples was proposed. This linear correlation between uptake and surface area was further examined and confirmed by Zuttel and co-workers (Figure 8.6).52,53 In another study, Kajiura et al.54 measured the hydrogen storage capacities of various commercially available carbon materials at room temperature and up to 8 MPa. The highest storage capacity was 0.43 wt.% for a purified SWNT. These low values are consistent with previously reported results 42 indicating that carbon materials are not suitable for hydrogen storage at room temperature. Poirier et al.55 investigated hydrogen adsorption on high specific surface area AC (activated carbon), CNFs and SWNTs with experimental and theoretical considerations. Some titanium free SWNTs appeared to surpass large surface area ACs in hydrogen uptake capacity both at room temperature and under cryogenic conditions. Hydrogen adsorption capacities of 0.2 wt.% and 4 wt.% were obtained for the SWNTs at 1 bar and room temperature and 1 bar and 77 K, respectively. It was proposed that due to the specific bundle structure, the SWNTs could favor hydrogen adsorption over the layered graphitic structures typically found in ACs. Schimmel et al.56,57 investigated the hydrogen storage capacities of activated charcoal, CNFs and SWNTs. They found that the hydrogen adsorption capacities of these materials correlated with their corresponding surface areas. The activated charcoal with a surface area of 2200 m2/g showed the highest hydrogen adsorption of 2.2 wt.% at 1 bar and 77 K. The adsorption potential of the activated charcoal was found to be 5 kJ/mol, which is in agreement with the result of Benard and co-workers.58 Hirscher et al.59,60 investigated the hydrogen storage capacity of different carbon nanostructures having

Figure 8.6 Desorbed amount of hydrogen versus the BET surface area (round markers) for carbon nanotube samples together with the ﬁtted line. Data from Nijkamp et al.51 (triangular markers) are also shown.53

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specific surface areas ranging from 22 to 2560m2/g. An almost linear relation between the storage capacity and the specific surface area was obtained at both room temperature and 77 K with the highest storage capacities of 0.5 and 4.5 wt.%, respectively (Figure 8.7). Kadono et al.61 examined the hydrogen adsorption properties for two different types of carbon nanomaterials at 77 K and pressure up to 2 MPa. They found that the material with subnanopores, which had a narrow distribution of pore size at approximately 0.5 nm diameter (Figure 8.8), showed a large hydrogen uptake of 2.8 wt.%, compared with 0.7 wt.% for the sample with a large pore diameter of 3 nm. Moreover, the estimated density of adsorbed hydrogen in pores having 0.5 nm diameter was found to be comparable to that of the density of bulk liquid hydrogen. Strobel et al.62 measured hydrogen adsorption on different carbon materials at 12.5 MPa and 296 K with BET (Brunauer-Emmet-Teller) surface areas ranging from 100 up to 3300 m2/g using a microbalance. The maximum adsorption of 1.6 wt.% was found for an activated carbon with a specific surface area of 1400 m2/g, but not on the material with the maximum surface area in the study. This indicated that only certain pores contributed to the hydrogen storage under that test condition. Zuttel et al.63 presented an empirical model for the hydrogen adsorption on carbon nanostructures. Hydrogen condensation as a monolayer at the surface of nanotubes and bulk

Figure 8.7 Hydrogen storage capacity of various carbon nanostructures versus the speciﬁc surface area at RT and at 77 K. The slopes of the curves are 0.23 103 and 1.91 103 wt.% g m2, respectively.60

Hydrogen Adsorption and Storage

Figure 8.8 Micropore size distribution of activated carbon ﬁbers obtained from the nitrogen adsorption isotherm at 77 K using the Horvath and Kawazoe method. Vp is the volume in the region for pores of less than 0.56 nm diameter, which is regarded as the effective pore diameter for hydrogen adsorption.61

condensation in the cavity of the tube was assumed in this model. In order to calculate the quantity of adsorbate in the monolayer they used the density of the liquid adsorbate and computed the volume of the molecule as Eqn. 8.11. VAd ¼

MAd rAd NA

Eqn. 8.11

where MAd represents for the molecular mass of the adsorbate, rAd for the density of the liquid adsorbate and NA for the Avogadro constant (NA ¼ 6.02201023 mol1). VAd is the volume reserved for each molecule of liquid. With the assumption that the molecules are spherical and closed-packed, the volume of the sphere representing the molecule is pﬃﬃﬃ by a factor of p/(3 2) smaller than VAd (Eqn. 8.12). p p MAd VM ¼ pﬃﬃﬃVAd ¼ pﬃﬃﬃ 3 2 3 2 rAd NA

Eqn. 8.12

From the volume of the spherical molecule the diameter dM can be calculated and assuming a closed package of the molecules in a two-dimensional layer at the surface, each molecule occupies the surface area SM (Eqn. 8.13 and Eqn. 8.14). ﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃ sp rﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃ ﬃﬃﬃ 3 2MAd 3 6VM dM ¼ ¼ Eqn. 8.13 p rAd NA pﬃﬃﬃ pﬃﬃﬃpﬃﬃﬃ 2 3 3 2MAd 3 2 ðdM Þ ¼ Eqn. 8.14 SM ¼ 2 2 rAd NA

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Table 8.1 The properties of hydrogen and nitrogen as adsorbates. M the molecular mass, r the density of the liquid at the boiling point, VM the volume of the molecule, d the diameter of the molecule, SM the surface area occupied by the molecule, 1/S the amount of adsorbate per surface area unit in a monolayer. Property N2 H2

M [g mol1] r [g cm3] VM [nm3] d [nm] SM [nm2] 1/S [mol m2] 1/S [g m2]

28.0140 0.8070 0.0426 0.4335 0.1627 1.02 105 2.86 104

2.0159 0.0708 0.0350 0.4059 0.1427 1.16 105 2.35 105

The above parameters for hydrogen and nitrogen are summarized in Table 8.1. The condensation of a monolayer of hydrogen on a graphene sheet with a specific surface area of S1 ¼ 1315 m2 g1 leads to m(H2)/m(C) ¼ 0.03085 according to Eqn. 8.15. mðH2 Þ S1 MAd ¼ SM NA mðCÞ

Eqn. 8.15

This converts to a maximum concentration of 2.28 103 mass% S [m2 g1] ¼ 3.0 mass% (H/C ¼ 0.18) hydrogen on carbon according to Eqn. 8.16 C½mass% ¼

1 mðCÞ

1 þ mðH2 Þ

Eqn. 8.16

After obtaining the maximum amount of hydrogen adsorption on the monolayer surface; they continued to calculate the hydrogen uptake in the cavity of the nanotube. As mentioned above, they assumed the condensation of the hydrogen gas in the cavity of a nanotube and thus the ratio of the mass of hydrogen to the mass of carbon can be expressed by Eqn. 8.17. mðH2 Þ S1 dNT rAd ¼ mðCÞ 4

Eqn. 8.17

From Eqn. 8.17, the ratio for single-wall nanotubes with a diameter of dNT ¼ 0.671 nm is m(H2)/m(C) ¼ 0.0156, which results in c ¼ 1.54 mass% (H/C ¼ 0.09). However, the available volume inside a nanotube is much smaller than the estimated volume in the above model. This is due to the space occupied by the carbon atoms with a covalent radius of approximately 0.077 nm. The calculated amounts of hydrogen absorbed at the surface and in the cavity of the carbon nanotubes are shown in Figure 8.9. From Figure 8.9, the surface absorption is highest for the CNT with the largest specific surface area, i.e. the SWNT. The bulk absorption in the cavity is proportional to

Hydrogen Adsorption and Storage

Figure 8.9 Calculated amount of adsorbed hydrogen on carbon nanotubes assuming condensation of hydrogen. (a) Monolayer adsorbed at the surface of nanotubes as a function of the number of shells (markers and line, axis: left and bottom). (b) Hydrogen condensed in the cavity of nanotubes as a function of the diameter (dashed lines, axis: left and top) for various numbers of shells, NS ¼ 1 (SWNT), NS ¼ 2, NS ¼ 5, NS ¼ 10 (dotted line, axis: left and top).63

the diameter of the tubes and is also highest for SWNT. Furthermore, the mass% of adsorbed hydrogen in the cavity of the tube increases with the tube diameter. The SWNT with a diameter of 2.2 nm has the potential to adsorb 5 mass% of hydrogen. However, this model suffers from the assumption that the hydrogen condenses in the nanotubes although the critical temperature of hydrogen is 33 K. The experimental hydrogen storage data correlated with the BET specific surface area and, the slope of the linear relationship is 1.5103 mass% S [m2 g1]. Therefore, the extrapolated maximum hydrogen capacity of a carbon sample is 2.0 mass%. The deviation between the experimental and calculated maximum hydrogen uptake was possibly due to the imperfection of the model. Furthermore, it can be concluded that the hydrogen sorption mechanism is related to the surface of the sample, i.e. a surface adsorption process. Hydrogen Storage in Carbon Nanotubes Carbon nanotubes can be envisaged as a cylindrical hollow tube rolled up from a single layer of graphene, diameter in the nanometer range, and length usually on the micron scale (Figure 8.10).9 The CNT is composed of only one sheet of graphene defined as single-wall nanotube. Since a graphene sheet can be formed into a tube in different ways, different types of carbon nanotubes exist, namely ‘armchair’, ‘zigzag’, and ‘chiral’ types (Figure 8.11).64

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Figure 8.10 Schematic of the structure of a carbon nanotube.9

Besides the single-walled nanotubes, multiple sheets can be rolled to form multiwall nanotubes (MWNTs). The MWNT is an arrangement of coaxial tubes of graphene sheets forming a tube-like structure. Each MWNT has from two to fifty such tubes, with inner diameters from 1.5 to 15 nm and outer diameters from 2.5 to 30 nm. The interlayer distance of the MWNT is 0.34e0.36 nm, similar to the interlayer distance in graphite. A schematic representation of a typical MWNT is shown in Figure 8.12.65 After the pioneering works of Dillon and co-workers,31 attempts were made to use these materials for hydrogen storage. Both SWNTs and MWNTs have been investigated. Chen et al.66 studied the hydrogen storage of aligned MWNTs synthesized by

Figure 8.11 Schematic of the different single-walled nanotube (SWNT) structures: (a) armchair, (b) zigzag, and (c) chiral.64

Hydrogen Adsorption and Storage

Figure 8.12 Schematic of the MWNT.65

a plasma-assisted CVD, or chemical vapor deposition, method. It was found that a hydrogen storage capacity of 5e7 wt.% was obtained at room temperature at 10 atm pressure for the as-synthesized sample. The hydrogen storage capacity increased up to 13 wt.% for a sample having removed catalyst tips by acid wash. However, the release of adsorbed hydrogen required heating to 300 C. The authors attributed their high hydrogen uptake value to the open end of the nanotubes and their alignment. Cao et al.67compared the hydrogen adsorption behaviors of well-aligned MWNTs and randomly ordered MWNTs at 290 K and 10 MPa. The hydrogen uptake of aligned MWNTs was much higher than that of randomly ordered MWNTs, with values of 2.4 wt.% and 0.5 wt.%, respectively (Figure 8.13). The authors proposed that intertube channels of the well-aligned MWNTs were the effective domains for the hydrogen adsorption. In another work, Wang et al.68 reported the hydrogen storage in CNTs and CNT films at ambient temperature and pressure. The CNT films adsorbed much more H2 than CNT powder, and a H2 adsorption capacity of 8 wt.% was achieved for the CNTs films. However, the mechanism of this hydrogen uptake was not clear. Hou et al.69 studied the hydrogen storage capacities for the as-prepared and purified MWNTs (Figure 8.14). The hydrogen uptake of the purified sample was much higher than that of the asprepared one. The hydrogen uptake of 6.3 wt.% and 31.6 kg/m3 was obtained for the purified sample at room temperature and 14.8 MPa. The authors attributed this to the opened tips, simpler chemical state, and larger micropore volume of the purified MWNT sample. Later, this group70 studied the hydrogen storage capacity of MWNTs with diameters ranging from 13 to 53 nm. They found that the hydrogen storage capacity of the MWNTs is proportional to their diameter. However, hydrogen in all

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Figure 8.13 Hydrogen adsorption of the aligned and random MWNTs.67

types of MWNTs measured could not be completely desorbed at room temperature and ambient pressure. The authors believed this undesorbed portion of hydrogen possibly corresponded to the chemisorption of hydrogen in MWNTs. To make the adsorption mechanism clear, they performed high-resolution transmission electron microscopy (HRTEM) observations on the MWNT samples. From the HRTEM images, a kind of discontinuous graphene stack structure defined as a ‘carbon island’ was found (Figure 8.15a). They proposed that these ‘carbon islands’ together with the defects or cavities around them could be the efficient hydrogen adsorption sites in MWNTs. A further heat treatment of MWNTs at 2573 K results in these ‘carbon islands’ (Figure 8.15b) disappearing. This disappearance is due to the sintering effect. In this case the treated samples had a much lower hydrogen storage capacity than the untreated ones.

(a)

(b)

Figure 8.14 TEM images of (a) the as-synthesized MWNTs and (b) the puriﬁed MWNTs.69

Hydrogen Adsorption and Storage

(a)

(b)

Figure 8.15 HRTEM images of MWNTs before (a) and after heat treatment (b).70

This result indicates indirectly that ‘carbon islands’ provide major hydrogen adsorption sites for MWNTs, although further investigation is needed. Shiraishi et al.71,72 studied the hydrogen storage properties for the SWNT bundles and peapods, i.e. C60 encapsulated SWNTs. The hydrogen desorption occurred at about 350 K through the TPD method. From the results it can be concluded that the inter-tube sites in SWNT bundles can be used as a host for hydrogen storage, and the availability of ‘sub-nanometer’-sized spaces was responsible for the moderate desorption temperature. However, a low value of H2 uptake of 0.3 wt.% was observed at room temperature and 9 MPa on SWNTs. Gundiah et al.73 investigated the hydrogen storage performance of several kinds of SWNT and MWNT samples obtained by different methods of synthesis, or subjected to different pre-treatment procedures. A maximum capacity of 3.7 wt.% was obtained for the acid-treated aligned MWNT sample at 300 K and 145 bar, which coincides well with the result obtained by electrochemical measurements. Gao et al.74 made a comparative study of the hydrogen storage properties of open-tipped MWNTs (OCCNT), high specific surface area activated carbon AX-21, close-tipped MWNTs (CTCNT), and SWNTs both at room temperature and at 77 K (Figure 8.16). Figure 8.16 shows the hydrogen adsorption isotherms for the four samples. From the figure, it is clearly shown that OTCNT possessed the highest hydrogen uptake capacity, with the value of 6.46 wt.%, among these materials at 77K. The hydrogen adsorption experiments at room temperature also indicated that the highest storage capacity of 1.12 wt.% was obtained for OTCNT. The high hydrogen storage capacity for OTCNT could be attributed to its high volume of subnanosized pores (<0.6 nm) which facilitate the micropore filling process. In another work, Jang et al.75 examined the hydrogen storage capacities of bamboo-shaped MWNTs, and compared them with those of SWNTs and MWNTs. It was found that the greatest hydrogen adsorption is about 0.4 wt.% at atmospheric pressure and temperature for the bamboo-shaped MWNT sample with

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Figure 8.16 Isotherms of wt.% H2 uptake vs. pressure at 77 K for AX-21, OTCNTs, SWNTs, and CTCNTs.74

a herringbonelike structure (Figure 8.17). The SWNTs showed a comparable hydrogen storage capacity, whereas the MWNTs were insensitive to the pressure change. Anson et al.76 studied the hydrogen uptake of SWNTs before (raw) and after heat treatment (HT) in air by measuring their adsorption isotherms under nearly 1 bar at

Figure 8.17 (a) TEM image of the BS-MWNTs synthesized by the vapor-phase growth. (b) Highresolution TEM image manifesting the open-ended (discontinuous) graphene sheets and herringbonelike structure of a BS-MWNT compartment.75

Hydrogen Adsorption and Storage

Figure 8.18 Experimental hydrogen adsorption isotherms at 77, 87 and 298 K for (a) raw sample and (b) HT sample.76

different temperatures. Hydrogen adsorption isotherms at 77, 87 and 298 K for raw and HT samples are shown in Figure 8.18a and b, respectively. From the data, it is easily seen that the HT sample absorbs more hydrogen than the raw material; although the maximum hydrogen uptake is only 1 wt.% at 77 K and the adsorbed amount at 298 K is very low. The experimental hydrogen isotherms were described and analyzed using the following form of a virial equation: m n X 1X a i yi þ bi yi Eqn. 8.18 ln p ¼ ln y þ T i¼0 i¼0 v ¼ amount adsorbed p ¼ pressure T ¼ temperature ai, bi ¼ empirical parameters Here Eqn. 8.18 is used to calculate the isosteric heats of hydrogen adsorption on SWNT samples. The isosteric heat of adsorption, Qst, is obtained from adsorption isotherms measured at different temperatures and is often used to characterize adsorption energies. Applying the Claussius-Clapeyron equation ln p ¼

Qst 1 þC R T

Eqn. 8.19

to Eqn. 8.18 leads directly to the following expression for Qst: Qst ¼ R

m X i¼0

ai vi

Eqn. 8.20

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Figure 8.19 Isosteric heats of hydrogen adsorption (Qst) on raw and HT samples calculated from Eqn. 8.20.76

where R is the universe gas constant. Qst calculated using Eqn. 8.20 and the fitted parameters is shown in Figure 8.19 for both samples. Error bars in the plots represent standard deviations of the Qst values estimated from variances of parameters calculated from the data fitting procedure. From Figure 8.19, the Qst decreases with hydrogen coverage. This is consistent with the fact that the binding energy for hydrogen in the pores is heterogeneous, and hydrogen molecules first occupy a lower energy configuration leading to larger isosteric heats of adsorption at lower loading. The Qst for the HT sample was higher than for the raw sample at any coverage, indicating that heat treatment increased the energy for H2 adsorption on SWNT. At zero coverage, the isosteric heats of adsorption were calculated to be 7.42 KJ/mol for the raw and 7.75 KJ/mol for the AT SWNTs. Lawrence and Xu77 reported that the hydrogen adsorption in SWNTs seemed to achieve saturation at 300 atm at room temperature and a capacity of 0.91 wt.% was obtained. This is lower than that of high specific surface area ACs. Taking into account the specific surface area of the SWNTs (800 m2/g) and ACs (2800 m2/g), adsorption rates of 0.53103 wt.% / 1000 m2/g and 1.14103 wt.% /1000 m2/g were obtained for the ACs and SWNTs, respectively. The latter value compared well with previous results.52,63 Chen and Huang78 reported the hydrogen storage properties of KOH modified MWNT. The high hydrogen adsorption value of 4.47 wt.% was obtained by TDS (temperature desorption spectroscopy) experiments, compared with 0.71 wt.% for the unmodified sample. The defect structure of CNTs formed via KOH activation was mainly attributed to this high hydrogen sorption capacity (Figure 8.20).

Hydrogen Adsorption and Storage

Figure 8.20 TEM micrographs (a) as-prepared MWNTs samples (b) KOH modiﬁed MWNTs sample.78

Zacharia et al.79 investigated the dependency of the gravimetric hydrogen storage capacity of SWNTs and MWNTs with the specific surface area at room temperature. The capacity scaled linearly with the specific surface area, irrespective of the type of nanotubes used at an equilibrium pressure of 1.6 MPa (Figure 8.21). A small amount of hydrogen uptake of less than 0.16 wt.%, lower than those calculated by the empirical condensation model, was found for all the samples. Hydrogen Storage in Graphite/Carbon Nanoﬁbers Graphite/carbon nanofibers (GNF/CNF) can be viewed as an arrangement of graphene platelets in a parallel, perpendicular, or angled orientation with respect to the fiber axis.

Figure 8.21 The linear scaling of the gravimetric hydrogen storage capacity of carbon nanotubes with speciﬁc surface area. In the ﬁgure, trace (- ) which indicates the ideal hydrogen storage capacity obtained using the empirical condensation model is compared with the experimental data (C).79

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(a)

(b)

(c)

Figure 8.22 Different structures of carbon nanoﬁbers (CNFs): (a) deck of cards structure, (b) herringbone structure, and (c) tubular structure.80

The arrangement of the graphite platelets depends on the geometry of the catalyst; on which CNF structures are formed at high temperatures from a mixture of carboncontaining gases and hydrogen, or from hydrocarbons alone.9 The platelets are structured according to the catalyst shape over a large number of platelet layers. Different shapes of CNF are shown in Figure 8.22.80 The fibers are usually between 5 and 100 mm in length, and their diameter is between 5 and 100 nm. The BET specific surface area of these ordered crystalline solids is normally between 100 and 300 m2 g1, but can sometimes reach values up to 700 m2 g1. In the ‘herringbone’ structure (Figure 8.22b) the graphite layers are separated from one ˚ . The structure of the platelets provides many edges on another by a distance of 3.4 A a nanofiber, which represent a large surface, with many interaction sites for adsorbates.9 Stimulated by the extraordinary results of Chambers and co-workers,39,40 many research groups investigated the hydrogen storage properties of GNFs/CNFs. However, none of the groups repeated the high hydrogen storage results. Gupta and Srivastava81 synthesized and investigated the hydrogen adsorption properties of GNFs at room temperature and pressure of 120 atm. The Pressure-Composition-Temperature isotherm showed hydrogen storage capacity of around 10 wt.% for the GNF prepared from ethylene gas. However, Marella and Tomaselli82 studied the hydrogen storage capacities of CNFs which were grown by catalytic decomposition of ethylene in the presence of hydrogen with different catalysts and treatment procedures. A maximum hydrogen storage only of 1.38 wt.% can be obtained for a Pd-doped CNF at 77 bar and room temperature. Even worse, Blackman et al.83 investigated the hydrogen storage capacities of activated CNFs by physical and chemical methods, respectively and found that the BET surface area increased to 1000m2/g after KOH activation; however, the increase of the hydrogen adsorption capacities did not correspond directly with the increase in the surface areas. Hydrogen storage capacity of less than 0.4 wt.% was obtained for all the

Hydrogen Adsorption and Storage

samples at 10 MPa and 30 C. Browning et al.84 studied the hydrogen storage for activated CNFs and found the hydrogen adsorption capacities can be up to 6.5 wt.% at 12 MPa and ambient temperature. This high capacity of hydrogen storage can be explained by the physisorption on the external surface area and the hydrogen dissociation on graphite edge sites, which is a chemisorption process. Hong et al.85 investigated the hydrogen storage properties of GNFs prepared from catalytic graphitization of electrospun poly (vinylidene fluoride) nanofibers. The hydrogen uptake of 0.11e0.18 wt.% was observed at room temperature and 80 bar. They suggested that the ultramicropores (< 1 nm) were responsible for this hydrogen adsorption. Hydrogen storage in activated carbon Activated carbon (AC) is composed of all assemblies of defective graphene layers. The pore diameters are usually less than 1 nm, and form a specific surface area up to 3000 m2/g. This highly developed porosity makes activated carbon widely used for purification of water, of air, and separation of gas mixtures.86 More recently, AC has attracted considerable attention for hydrogen storage. In 1994, Chahine and Bose87 investigated the hydrogen storage capacity of a series of commercially available activated carbons at 77 K and 1 MPa. They found maximum hydrogen uptake of 5.3 wt.% was obtained for AX-21, with the BET surface area of 3000m2/g . Zhou et al.88 compared the hydrogen storage properties of activated carbon AX-21 and MWNT at ambient temperature. Values of 0.5 wt.% and 0.14 wt.% were obtained, respectively, at 298 K and 6 MPa. In addition, the corresponding hydrogen adsorption enthalpies for these two samples were calculated to be 6.4 and 1.7 kJ/mol. De La Casa-Lillo et al.89 reported hydrogen sorption capacities for a large variety of ACs and ACFs (activated carbon fibers). The highest value of hydrogen adsorption was found to

Figure 8.23 Hydrogen adsorption isotherms for a series of activated carbons and activated carbon ﬁbers at 293 K and up to 70 MPa.89

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Figure 8.24 H2 adsorption capacity versus (a) total surface area for 1 and 10 bar and (b) N2 and CO2 micropore volume for 10 bar.90

be 1 wt.% at room temperature and a pressure of 70 MPa (Figure 8.23). The experimental results reported correlated well with those of the theoretical studies12,14; and confirmed that the optimum pore size for hydrogen adsorption in porous carbons is about 0.6 nm, or sufficient to accommodate two layers of adsorbed hydrogen. Texier-Mandoki et al.90 examined hydrogen storage properties of different ACs with surface areas ranging from 890 to 3000m2/g (Figure 8.24). They found that hydrogen storage capacities correlated better with micropore volume measured by CO2 adsorption than by BET surface area. They proposed that the ultramicropores, pores smaller than 0.7 nm, played an essential role in the hydrogen uptake process. This is in agreement with the work by Zhao and co-workers.91 A maximum hydrogen adsorption of 4.5 wt.% was obtained for an AC sample having a surface area of 3000 m2/g, and micropore volume of 0.6 cm3/g at 77 K and 10 bar. Hydrogen Storage in Other Carbon Nanostructures Other carbon nanostructures such as graphite, carbon aerogel, mesoporous carbon etc have also been investigated for hydrogen storage. Kabbour et al.92 presented the hydrogen sorption properties of activated (ACA) and metal-doped CA (carbon aerogel) (Figure 8.25a). A maximum hydrogen adsorption capacity of 5.3 wt.% was found for the ACA sample with a surface area of 3200m2/g at 77 K. A linear dependence of H2 excess adsorption with the BET surface area was also established (Figure 8.25b). The hydrogen enthalpy of adsorption for the metal-doped CA was measured to be about 7 kJ/mol, which is larger than that of AX-21.58 Orimo et al.93,94 showed that nanostructured graphite can hold up to 7.4 wt.% hydrogen. The graphite must be prepared by reactive ball milling for 80 hours under a 1.0 MPa hydrogen atmosphere. However, hydrogen desorption began at about 600 and 950 K, respectively,94 where 80% of the total amount of hydrogen was desorbed at the first desorption temperature. In another work, Chen et al.95 investigated the

Hydrogen Adsorption and Storage

Figure 8.25 (a) Adsorption isotherms at 77 K for the ACAs. (b) Excess gravimetric density (H2 wt.%) saturation value at 77 K as a function of BET surface area.92

hydrogen adsorption/desorption behavior of mechanically milled graphites under different hydrogen pressures ranging from 0.3 to 6 MPa, and found the chemisorbed hydrogen concentration decreased with increased hydrogen pressure. However, the physisorbed hydrogen concentration increased with increased hydrogen pressure. The maximum uptake of hydrogen for chemisorption and physisorption was 6.1 and 0.5 wt.%, respectively. Pang et al.96 examined the hydrogen adsorption of mesoporous carbon materials with different surface areas, ranging from 886 to 2314 m2/g, and different pore sizes and volumes. They found that the maximum hydrogen adsorption capacity was 1.78 wt.% at 77 K and ambient pressure for the sample with the highest surface area. Furthermore, an approximate linear relationship between hydrogen adsorption capacity and surface areas of the samples was found (Figure 8.26). Shao et al.97 investigated the hydrogen storage behavior of an activated mesocarbon microbead sample with a BET surface area of 3180 m2/g both experimentally and theoretically. Hydrogen adsorptions of 0.3 wt.% at 298 K and 7.8 wt.% at 77 K under the ambient pressure were obtained experimentally. However, at a pressure of 10 MPa, the theoretical prediction showed much higher values of 3.2 and 15 wt.%, respectively. Gogotsi et al.98 reported the hydrogen storage properties of CDC (carbide-derived carbon), which has a specific surface area of up to 2000m2/g and open pore volume up to 80%. A maximum hydrogen adsorption capacity of 3 wt.% was obtained at 77 K and 1 bar. It was found that the small pores are responsible for the high sorption ability. This supports the theoretical prediction that small pores with narrow size distribution are the key to high hydrogen uptake. Yang et al.99 reported hydrogen sorption properties of zeolite-like carbon materials with surface area up to 3200 m2/g (Figure 8.27a). A maximum hydrogen adsorption of 6.9 wt.% was achieved at 77 K and 20 bar for a sample with a high proportion of micropore surface area and micropore volume in addition to

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Figure 8.26 The correlation btween hydrogen adsorption capacity and the speciﬁc surface area of the mesoporous carbon materials: OC, FC, DC-1,-2, and HC at 77 K and 0.11 MPa.96

a high BET surface area. This capacity was among the highest values ever reported for carbon or any other porous material, including MOFs (metal organic frameworks) or zeolites. Moreover, the hydrogen enthalpy of this sample was calculated to be 8.2 kJ/mol K, which was also among the highest values reported for carbon materials. This suggests a strong interaction between adsorbed hydrogen and the carbon surface (Figure 8.27b). This high heat of adsorption and the “optimal” pore size 0.6e0.8 nm of this material contributed to the high hydrogen sorption ability. Terres et al.100 reported the H2 storage capacities of spherical nanoporous carbons with surface areas ranging from 946 to 1646 m2/g. The maximum hydrogen uptake of 2.7 wt.% was obtained at 77 K for a sample with a periodic array of channels and surface area of 1642 m2/g. The authors emphasized again the importance of surface area to the hydrogen storage capacity of carbonaceous materials.

Figure 8.27 (a) Hydrogen isotherms for a series of zeolite-like carbon materials. (b) Isosteric heat of H2 adsorption (Qst) for sample d in (a) as a function of the amount of hydrogen adsorbed.99

Hydrogen Adsorption and Storage

Hydrogen Storage in Carbon Nanostructures by Spillover Hydrogen spillover is defined as the dissociative chemisorption of hydrogen on the catalyst; and the subsequent migration of atomic hydrogen onto the surface of the receptors such as activated carbon. This new technique will open a new direction of finding high performance hydrogen storage materials. Yang et al.101e104 investigated the hydrogen storage properties in various carbon nanostructures by hydrogen spillover. The simple mixing of CNTs with supported palladium was found to increase the uptake of the hydrogen by a factor of three.103 In another study, Mu et al.105 studied the hydrogen storage properties of modified CNTs and compared them with those of original samples. It was found that hydrogen uptake improved 3-fold for the treated samples, which can be up to 4.5 wt.% at room temperature and 10 MPa. It is proposed that the wall defects or disordered structure brought about by microwave plasma etching and Pd decoration can contribute to this increased hydrogen storage performance. The mechanism of hydrogen uptake included the dissociation of hydrogen molecules into atoms or protons on the surface of metallic Pd, diffusion into the hollow interiors of tubes, and recombining with the hydrogen molecules. Zacharia et al.106 compared the hydrogen sorption capacity of platinum-MWNT with that of a pristine one and found the hydrogen storage capacity of the platinum deposited nanotube was nearly 40 times higher than that of the pristine sample, which can be up to 3.1 wt.% at 298 K and 16 atm (Figure 8.28). The enhanced hydrogen sorption capacity was attributed to the spillover of physisorbed hydrogen from platinum particles. Lee et al.107 reported the hydrogen storage properties of ACFs modified by Ni doping and fluorination. After treatments, the specific surface area of ACFs decreased, however,

Figure 8.28 Hydrogen storage capacity of Pt-doped MWCNTs measured at 298 K and under an equilibrium pressure of 1.67 MPa. Inset of the ﬁgure displays the hydrogen storage capacity of pristine MWCNTs measured under similar conditions.106

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the hydrogen uptake increased compared with those of unmodified sample. The effect of catalytic activation of Ni and enhancement of hydrogen affinity by fluorination can contribute to the higher hydrogen uptake capacity. Thus, they concluded that the hydrogen storage capacity was not determined only by the factor of specific surface areas, but it was also related to the micropore volumes, pore size distribution and surface properties of these ACFs. The above results indicate that the early spectacular hydrogen storage capacities reported for carbon nanostructured materials at near-ambient temperature have not been independently verified by other groups. In fact, repeatable hydrogen adsorption of approximately ~ 5 wt.% at 77 K and < 0.5 wt.% at room temperature can be achieved under high-pressure conditions for the carbon based nanostructured materials. The maximum hydrogen uptake correlates linearly with the specific surface area and pore volume of the adsorbent. This suggests that the best materials for practical storage application will be those with the largest specific surface per unit volume. Since the physisorption of hydrogen on carbon nanostructures falls short of the DOE targets for the use of hydrogen as a transportation fuel, further efforts at designing materials with high surface areas and pore volumes are needed.

Hydrogen Storage of Metal-Organic Frameworks Recently, a new class of porous metal-organic frameworks (MOFs)108e118 have attracted much attention for hydrogen storage because of their exceptionally high specific surface area and microporosity, uniform but tunable pore size, very low density and well-defined hydrogen occupation sites. Generally speaking, these materials are synthesized in a building block approach from metal or metal oxide vertices interconnected by organic linkers. They have the potential to exhibit properties inherent to the building blocks, such as geometric rigidity, chemical functionality, or chirality. A variety of linker molecules can be used to create an entire family of materials having different pore sizes and containing different chemical functionalities within the linkers, but all with the same basic framework topology. In addition, different metal corners and different linkage chemistries can yield a wide variety of other framework topologies. Yaghi and co-workers119 first reported the hydrogen adsorption properties of MOF-5 (Zn4O clusters linked by BDC; BDC ¼ 1,4-benzenedicarboxylates; see Figure 8.29) at both 78 K and ambient temperature. Promising hydrogen storage capacities of 4.5 wt.% at 78 K and ambient pressure and 1 wt.% at room temperature and a pressure of 20 bar were obtained. Moreover, they showed that the amount of adsorbed hydrogen increased with increasing pressure in a nearly linear fashion at both 77 and 298 K (Figure 8.30). The authors claimed that the observed sharp uptake of H2 at lower pressure indicates favorable sorption interactions between the MOF-5 framework and H2. They also proposed the presence of two

Hydrogen Adsorption and Storage

Figure 8.29 Single-crystal X-ray structures of MOF-5 illustrated for a single cube fragment of their respective cubic three-dimensional extended structure. On each of the corners is a cluster [OZn4(CO2)6] of an oxygen-centered Zn4 tetrahedron that is bridged by six carboxylates of an organic linker. The large spheres represent the largest sphere that would ﬁt in the cavities without touching the van der Waals atoms of the frameworks. Hydrogen atoms have been omitted.119

well-defined binding sites for hydrogen, the zinc and the BDC linker, through the measurement of inelastic neutron scattering. One year later, the same group120 examined the hydrogen storage capacities of a variety of MOFs (Figure 8.31) with surface areas ranging from 1466 to 4526 m2/g. The maximum hydrogen uptake in this research was 1.6 wt.% at 77 K and 1 bar for IRMOF-11, which had a surface area of 1911 m2/g. Apparently, there is a lack of correlation between the measured surface area and the hydrogen uptake. Moreover, the isotherms for hydrogen adsorption in different MOF materials did not achieve a plateau.

Figure 8.30 Hydrogen gas sorption isotherm for MOF-5 at (a) 78 K and (b) 298 K.119

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Figure 8.31 (a) Isoreticular (having the same underlying topology) MOF, Zn4O(L)3, are constructed by linking zinc oxide clusters with linear carboxylates L such as those shown. (b) The structure of MOF177, Zn4O(BTB)2, is formed by linking the same clusters with a trigonal carboxylate.120

This implies that no saturation took place at the surface (Figure 8.32). They also found that the MOF-5 (IRMOF-1) gave a hydrogen uptake of 1.3 wt.%, not 4.5 wt.% as in a previous study119 at 77 K and 1 bar. This is consistent with the result of Panella and coworkers.121 In a more recent investigation,122 Yaghi and co-workers reported the hydrogen sorption properties for a variety of MOFs with surface areas ranging from 1070 to 5040 m2/g at 77 K and hydrogen pressures up to 90 bar. They found that H2 saturation uptake took place between 25 and 80 bars for all the samples (Figure 8.33) and; more interestingly, the hydrogen saturation uptake in these samples correlated well with the surface area just as for zeolites123 and carbon materials.51, 59 The maximum hydrogen uptake was 7.5 wt.% for IRMOF-177, which had the highest surface area in the study. The hydrogen adsorption data were also compared on a volumetric basis since this is of particular interest for hydrogen storage in transportation applications. Structures with high surface area and moderate density provided the highest storage density on a volumetric basis of 34 kg/m3.

Hydrogen Adsorption and Storage

Figure 8.32 Hydrogen isotherms for the MOF materials measured gravimetrically at 77 K (adsorption, 120 C desorption, B).

Here we should note that MOF-5, in this study, shows a hydrogen adsorption capacity of 5.2 wt.%, which is different from the above reported results.119e121 The variation of hydrogen uptake in different groups can be attributed to different synthesis methods, incomplete evacuation of the pores and/or partial decomposition of the framework upon exposure to air. To resolve these discrepancies and determine the true hydrogen storage capacity of MOF-5, Kaye et al. explored a synthesis method with complete activation and protection of the sample from air and water and found that MOF-5 exhibited a maximum ever reported excess hydrogen uptake of 7.1 wt.% at 77 K and 40 bar and total hydrogen uptake of 11.5 wt.% at 77 K and 170 bar, corresponding to a volumetric storage density of 78 Kg/m3 (Figure 8.34).124

Figure 8.33 High-pressure H2 isotherms for a series of MOFs at 77 K in gravimetric units (mg/g); solid markers represent adsorption, open markers denote desorption.122

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Figure 8.34 Hydrogen adsorption isotherms for MOF-5 at both 77 and 298 K, shown as excess (squares) and total (circles) uptake.124

Inspired by the initial promising performance of the MOFs, numerous new MOFs have been synthesized and their hydrogen adsorption capacities reported. Pan et al.125 reported the hydrogen sorption properties of a novel MMOM (microporous metal 0 coordination material): [Cu(hfipbb)(H2hfipbb)0.5] [H2hfipbb ¼ 4,4 -(hexafluoroisopropyl idene)bis(benzoic acid)], and found the adsorption capacity of this sample was close to 1 wt.% at 48 atm and room temperature. They attributed this hydrogen sorption behavior to the small pore dimension of the new material. They also pointed out that an optimal material would have pores that match the size of the gas molecules and have the largest possible pore volume. Kesanli et al.126 examined the hydrogen sorption capacities of two novel fourfold interpenetrated MOFs, Zn4O (L1)3 and Zn4O (L2)3 (L1 ¼ 6,60 -dichloro-2,20 -diethoxy-1,10 -binaphthyl-4-40 -dibenzoate, L2 ¼ 6,60 dichloro-2,20 -dibenzyloxy-1,10 -binaphthyl-4,40 -dibenzoate), and found hydrogen adsorption of 1.12 and 0.98 wt.% at room temperature and 48 atm, respectively. The ˚ due to the interpenetration can attribute the high small pore size of less than 5 A hydrogen uptake for these materials. Dybtsev et al.127 reported the hydrogen sorption behavior for the microporous manganese formate (Mn(HCO2)2$1/3 dioxane) with a BET surface area of only 240 m2/g (Figure 8.35). The moderate hydrogen uptake of 0.9 wt.% was achieved at 77 K and 1 atm. This material showed selective sorption of H2 and CO2 over Ar, N2. Such selective sorption of H2 in this material is rare for microporous solids. The authors attributed this property to the small aperture of the channels of this MOF material. In a more recent investigation,128 they studied the hydrogen sorption properties of a highly porous MOF, Zn2(bdc)2(dabco)[bdc ¼ benzene-1,4-dicarboxylate, dabco ¼ 1,4-diazabicyclo[2.2.2]octane] with a surface area of 1450m2/g at 78 K and 1 atm. The maximum hydrogen uptake of 2.0 wt.% was attained under the test conditions.

Hydrogen Adsorption and Storage

Figure 8.35 X-ray crystal structure of Mn(HCO2)2 showing channels along the b-axis. Guest molecules are omitted. Each octahedron represents Mn(II) cation coordinated by six formate ligands.127

Furthermore, a rare guest-dependent dynamic behavior, in which the framework shrinks upon inclusion of organic guest molecules and expands upon guest removal, was revealed in this material. However, the reason for this is still not clear. Lee and Suh129 studied the hydrogen storage capacity of a robust porous metal-organic open framework, Ni(cyclam) (bpydc) [cyclam ¼ 1,4,8,11-tetraazacyclotetradecane, bpydc ¼ 2,20 -bipyridyl-5,50 dicarboxylate], with a surface area of 817 m2/g and pore volume of 0.37 cm3/g. It adsorbed H2 gas up to 1.1 wt.% at 77 K and 1 atm. Chen et al.130 reported the hydrogen adsorption behavior of a MOF with open metal sites, Cu2(bptc) [bptc ¼ biphenyl3,30 ,5,50 -tetracarboxylate] (MOF-505) at 77 K and 1 atm. The maximum sorption of 2.47 wt.% was gained for the fully activated sample (Figure 8.36). This value represents the highest hydrogen uptake under a pressure of 1 bar. It was proposed that that the open metal sites were responsible for this high adsorption capacity. Later, the same groups131 further investigated the hydrogen storage property for a new microporous dynamic MOF Zn(NDC)(4,40 -Bpe)0.5$2.25DMF$0.5H2O[NDC ¼ 2,6-naphthalenedicarboxylate, 4,40 -Bpe ¼ 4,40 -trans-bis(4-pyridyl)-ethylene]. This new material showed a maximum adsorption of 2.0 wt.% at 40 bar and 77 K. Considering its low surface area and pore volume, which are only 303 m2/g and 0.20 cm3/g, respectively, this hydrogen sorption amount was pretty surprising. The authors attributed this high hydrogen sorption capacity to the full usage of the pores in the MOF at higher pressure. This study also revealed the potential use of dynamic MOF materials for highpressure hydrogen storage; although their permanent porosity might be low because of framework transformations/deformations. Lee et al.132 reported the hydrogen adsorption of two microporous MOFs, [Zn3(bpdc)3bpy]$4DMF$H2O (1) and [Co3(bpdc)3bpy]$ 4DMF$H2O (2) (bpdc ¼ biphenyldicarboxylate; bpy ¼ 4,40 -bipyridine; DMF ¼ N,Ndimethylformamide) at both 87 and 77 K. It was found that the Co structure (2) adsorbed

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Figure 8.36 H2 isotherms for MOF-505 (77 K) after three different activation stages: stage I (25 C, squares), stage II (70 C, triangles), and stage III (120 C, circles); solid shapes indicate adsorption and open shapes indicate desorption data points.130

more hydrogen than the Zn structure (1) at both temperatures and all pressure levels. The values were 1.48 and 1.32 wt.% at 87 K and 1 atm, 1.98 and 1.74 wt.% at 77 K and 1 atm, respectively (Figure 8.37a). This difference was in agreement with the estimated total micropore volumes obtained from Ar adsorption data, which were 0.38 and 0.33 ml/g for 2 and 1, respectively. The isosteric heat of adsorption (Qst) of 1 and 2 were calculated using a variant of the Clausius-Clapeyron equation taking the hydrogen adsorption isotherm data at both 77 and 88 K:

Figure 8.37 (a) H2 adsorption (Ads) and desorption (Des) isotherms at 77 and 87 K for 1 and 2, respectively. (b) Isosteric heats of H2 adsorption for 1 and 2 as functions of the amount of hydrogen adsorbed.132

Hydrogen Adsorption and Storage

P1 T2 T 1 ¼ DHads ln P2 R T1 T2

Eqn. 8.21

where Pn ¼ pressure for isotherm n Tn ¼ temperature for the isotherm n R ¼ gas constant. Qst as a function the amount of hydrogen adsorbed amount for 1 and 2 is illustrated in Figure 8.37b. From these results, it is clearly shown that Qst values are higher for 2 than for 1 over the weight range 0.1e1.3%, indicative of somewhat stronger sorbentesorbate interactions, and, thus, higher energy of hydrogen adsorption, in 2. The Qst values of both 1 and 2 decrease as the hydrogen uptake increases. This is usually observed for MOFs. Chun et al.133 reported the hydrogen adsorption behavior of a series of isomorphous Zn MOFs that have a three-dimensional connectivity and interconnected pores. The surface areas of this series of samples ranged from 1450 to 2090 m2/g. The corresponding hydrogen sorption capacities of 1.7e2.1 wt.% were obtained at 77 K and 1 atm. Through detailed analysis of hydrogen sorption data in these MOFs, they also concluded that smaller pores actually took up hydrogen more effectively than very large ones; agreeing with previously reported results by Chen and co-workers.130 In 2004, Zhao et al.134 reported the hydrogen adsorption properties for Ni-based porous MOFs, Ni2(bipy)3(NO3)4 (bipy ¼ 4,40 -bipyridyl) and Ni3(btc)2(3-pic)6(pd)3 (btc ¼ 1,3,5-benzene tricarboxylate, pic ¼ 3-picoline, and pd ¼ 1,2-propanediol). The maximum hydrogen sorption capacity of 2.1 wt.% was achieved at 77 K and 1 atm. In the case of hydrogen adsorption on Ni2(bipy)3(NO3)4, a remarkable hysteresis for hydrogen desorption was observed due to the dynamic structure of the material (Figure 8.38). Hysteresis decreasing at higher temperature was explained by kinetic trapping of the hydrogen. This kinetic trapping effect is related to the flexibility of the porous framework materials. Kaye et al.135 reported the hydrogen adsorption properties for dehydrated Prussian blue analogues M3[Co(CN)6]2 (M ¼ Mn, Fe, Co, Ni, Cu, Zn). Hydrogen sorption capacities of 1.4 to 1.8 wt.% were obtained at 77 K and 1 atm. The hydrogen binding enthalpies for this series of cyano-bridged frameworks were 5.9e7.4 kJ /mol, higher than those of MOF-5120 and AX-21.58 At the same time, Chapman et al.136 examined the hydrogen sorption behavior for a series of Prussian blue analogues M3[Co(CN)6]2 (M ¼ Mn, Fe, Co, Ni, Cu, Zn, Cd). Maximum hydrogen sorption of 1.2 wt.% was achieved at 77 K and 1 atm. This compares well with the results of Kaye and Long.135 The hydrogen sorption capacities were found to depend on the different coordinatively unsaturated divalent metal cations contained in the materials. Later on, Long’s group137 reported the hydrogen sorption properties of a Mg-based MOF, which had very small pores of about 3.5 A˚ in diameter. The hydrogen uptake ability for this sample was only 0.6 wt.% at 77 K; however, the heat of adsorption was as high as 9.5 kJ/mol, which is

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Figure 8.38 Adsoprtionedesorption isotherms for H2 adsorption (solid symbols) and desorption (open symbols) on adsorbents at 77K (A) activated carbon (squares) and Ni3(btc)2(3-pic)6(pd)3 (triangles); (B) Ni2(bipy)3(NO3)4 with different structural topologies (circles and inverted triangles).134

higher than that of previously reported Zn-based MOFs.120 This higher binding affinity can be attributed to the increased van der Waals contact, which is associated with the very small pore size. More recently, they138 reported the hydrogen storage properties of the dehydrated Prussian blue type solids Ga[Co(CN)6]2, Fe4[Fe(CN)6]3, M2[Fe(CN)6] (M ¼ Mn, Co, Ni, Cu) and Co3[Co(CN)5]2 and compared them with those of M3[Co (CN)6]2 (M ¼ Mn, Fe, Co, Ni, Cu, Zn). The maximum saturation hydrogen uptake of 2.3 wt.% was predicted at 77 K for Cu2[Fe(CN)6]. The hydrogen enthalpies for these samples were calculated to be 6e7 kJ/mol, which is consistent with the previous results for dehydrated Prussian blue analogues.135 Based on the fact that the hydrogen enthalpies were not substantially higher for the framework containing coordinatively unsaturated metal centers, they drew the conclusion that the exposed metal coordination sites in these materials did not dominate the hydrogen binding interaction. In another work, however, Yang and Zhong139 performed a systematic molecular simulation study on the adsorption and diffusion of hydrogen in MOFs. It was found that the metal-oxygen clusters were preferential adsorption sites for hydrogen in MOFs. The effect of organic

Hydrogen Adsorption and Storage

linkers became evident with increasing pressure, which was consistent with preciously reported results.119 It was also pointed out that MOFs were promising hydrogen storage materials at 77 K, but the hydrogen storage capacity at room temperature and moderate pressure was still short of expectations. Later, this same group140 conducted a theoretical investigation of the hydrogen adsorption in MOF-505, a MOF with open metal sites, by a combined GCMC simulation and DFT calculation. The metal-oxygen clusters were preferential adsorption sites for hydrogen, and the strongest adsorption of the hydrogen was found in the directions of coordinately unsaturated open metal sites. This provides evidence that the open metal sites have a favorable impact on the hydrogen sorption capacity of MOFs. However, the storage capacity of MOF-505 at room temperature and moderate pressure was predicted to be low. Krawiec et al.141 reported the hydrogen storage capacity for Cu3(BTC)2 (BTC ¼ 1,3,5-benzene tricarboxylate). The maximum hydrogen uptake of 2.18 wt.% was achieved at 77 K and 1 atm. This value is higher than the results of Lee and co-workers,142 which showed a hydrogen storage capacity of 1.44 wt.% at 77 K and 1 atm. The authors ascribed this improvement to the higher pore volume of the framework achieved by employing a novel synthesis procedure. The hydrogen binding enthalpy was calculated to be 6.1 kJ/mol, which compares well with values determined by Kaye and co-workers135 for dehydrated Prussian blue analogue materials. Also, Panella et al.143 studied the hydrogen adsorption in two different MOFs, MOF-5 and Cu-BTC at temperatures ranging from 77 K to room temperature. Heats of adsorption of 3.8 and 4.5 kJ/mol were obtained for MOF-5 and Cu-BTC, respectively. At room temperature both samples showed hydrogen storage capacities less than 0.4 wt.%, and at 77 K the saturation hydrogen uptake was 3.6 wt.% for Cu-BTC and 5.1 wt.% for MOF-5 (Figure 8.39). This is in good agreement with Yaghi’s result.122 A nearly linear relation between the maximum hydrogen uptake and surface area was found. They claimed that the low-pressure region of the adsorption isotherm was strongly dependent on the strength of the interaction between hydrogen molecules and the adsorbent; therefore, the difference in the storage capacity at low pressures for different materials was attributed to the difference in heats of adsorption. However, the saturation value was independent of the different heat of adsorption values, and could be correlated to the specific surface area of the adsorbing materials. Since 2003, Ferey and co-workers have reported the hydrogen sorption properties of several “MIL” series of MOFs.144e147 In 2003, they first reported the hydrogen adsorption capacities for two nanoporous MIL-53, Al(OH)(BDC) and Cr(OH)(BDC) at 77 K.144 The chromium compound showed a maximum adsorption of 3.1 wt.% under 16 bar, whereas the aluminum compound exhibited 3.8 wt.% at the same pressure. However, hysteresis in the adsorption and desorption was found, and adsorption capacities of 1.8 and 2.2 wt.% were recoverable from the Cr- and the Al-based samples, respectively, by desorption down to 0.1 atm (Figure 8.40). The mechanism of this phenomenon was not clear.

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Figure 8.39 D2 sites in Cu3(BTC)2: axial Cu site (left); view along [111] in the 5 Å small pore with 3.5 Å side windows (middle); and view along [100] showing the 9 Å pore.140

In a later study145 they synthesized a MOF MIL-101 with stoichiometry Cr3F (H2O)2O[(O2C)C6H4(CO2)]3$nH2O. This material had a zeotype cubic structure with a very large cell volume (~702,000 A˚3), and hierarchy of very large pore sizes ˚ ) with a nitrogen pore volume of 2.0 cm3/g and Langmuir (approximately 30e34 A surface area of ~5900 m2/g (Figure 8.41). Hydrogen adsorption studies showed that the material adsorbed 4.5 wt.% at 77 K and 3 MPa.

Figure 8.40 PCI curves (:: adsorption; ;: desorption) measured at 77 K up to 1.6 MPa of hydrogen pressure for the MIL-53 samples with M ¼ Cr3þ (open symbol) and M ¼ Al3þ (full symbol).144

Hydrogen Adsorption and Storage

Figure 8.41 (A) The computationally designed trimeric building block chelated by three carboxylic functions. The super tetrahedron was constructed with (B) terephthalic acid, which lies (C) on the edges of the super tetrahedron. (D) Ball-and-stick representation of one unit cell, highlighting one super tetrahedron drawn in a polyhedron mode. (E) Schematic 3D representation of the MTN zeotype architecture (the vertices represent the centers of each super tetrahedron) with the medium (20 tetrahedra) and large (28 tetrahedra) cages delimited by the vertex sharing of the ST.145

Later, hydrogen storage properties for two chromium carboxylates, MIL-100 and MIL-101, were investigated comprehensively.146 The maximum hydrogen adsorption capacity of 6.1 wt.% was obtained for the MIL-101 with a surface area of approximately 5500 m2/g at 77 K and 40 atm. The hydrogen enthalpy of this sample was calculated to be 10 kJ/mol. The authors attributed this high hydrogen uptake behavior to the small pore size and large micropore volume, combined with the coordinately unsaturation Cr metal centers. More recently147 they synthesized and investigated the hydrogen storage properties for a new three-dimensional chromium (III) naphthalene tetracarboxylate, Cr3O(H2O)2Ff1.5$ 6H2O (MIL-102). This new porous chromium carboxylate material exhibited a hydrogen storage capacity of 1.0 wt.% at 77 K and 35 bar. This low hydrogen sorption was attributed to the small surface area, only 42 m2/g, and the small free diameter of tunnels. In 2006, Sun et al.148 prepared and examined the hydrogen sorption capacity of a novel MOF Cu3(TATB)2(H2O)3 (TATB ¼ 4,40 400 -s-triazine-2,4,6triyltribenzoate) with a surface area of 3800 m2/g. Hydrogen uptake of 1.9 wt.% was obtained at 77 K and 1 bar. The authors ascribed this hydrogen uptake to the specific design factors of this material. The presence of accessible unsaturated metal centers and the existence of pores and channels in a size range well-suited to the H2 molecule. Later,

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(a)

(b)

Figure 8.42 (a) Crystal structure of PCN-9 showing an octahedral cage; the sphere represents void inside the cage. (b) View of the crystal structure of PCN-9 along the a-axis; the two interpenetrated networks are also shown.149

the same group149 reported the hydrogen sorption properties of a MOF with entatic metal centers, PCN-9 (Figure 8.42). This material exhibited a 1.53 wt.% hydrogen uptake at 77 K and 1 atm. The calculated heat of adsorption of 10.1 kJ/mol was among the highest yet observed for MOF materials.120,135,137This strong affinity was attributed to the existence of entatic metal centers in the sample. Dinca et al.150 reported hydrogen storage properties for a series of microporous MOFs constructed by the 1,4-benzeneditetrazolate ligand. It was found that the maximum adsorption capacity of 1.46 wt.% was achieved at 77 K and 1 atm for Zn3(BDT)3 [BDT2- ¼ 1,4-benzeneditetrazolate] with a surface area of 640 m2/g, the highest in this study. Furthermore, the enthalpies of adsorption in the range 6.0e 8.8 kJ/mol were calculated for these materials, higher than previously reported results for MOF-5120 and dehydrated Prussian Blue analogues.135 More recently, the same group151 reported the hydrogen storage capacity of microporous MOFs with exposed Mn2þ coordination sites. Total hydrogen storage capacity of 6.9 wt.% was obtained at 77 K and 90 bar together with the volumetric storage capacity of 60 kg/m3. In addition, a maximum enthalpy of adsorption of 10.1 kJ/mol was achieved. The H2 binding at coordinately unsaturated Mn2þ centers was attributed to the high hydrogen storage capacity and enthalpy. This ability of H2 to bind to metal atoms allows the H2 molecules to pack more closely together, and is expected to provide a major boost in storage capacity over simple H2 adsorption at nonmetal sites. Forster et al.152 studied hydrogen adsorption on nanoporous nickel (II) phosphates VSB-1 and VSB-5 (Figure 8.43). It was found that VSB-5 adsorbed substantially more hydrogen at low pressure. TPD studies showed that hydrogen had a strong interaction with VSB-5 with the main desorption peak at 109 K and a small shoulder at 149 K. The strong interaction was

Hydrogen Adsorption and Storage

(a)

(b)

Figure 8.43 Representations of VSB-1 (a) and VSB-5 (b) viewed down the c-axis, with octahedral nickel and tetrahedral phosphorus.152

confirmed by inelastic neutron scattering studies, whereas VSB-1 did not show a similar interaction. It was proposed that coordinatively unsaturated Ni2þ sites accessible to H2 molecules existed in VSB-5 resulting in this unique surface chemistry. Later on, they153 prepared a nickel sulfoisophthalate MOF NaNi3(OH)(SIP)2 [SIP ¼ 5-sulfoisophthalate], the hydrogen sorption properties of which were investigated at 77 K. The hydrogen uptake of this material was 0.94 wt.%. The strongest hydrogen binding site was found to be the coordinately unsaturated metal sites using neutron scattering spectroscopy. A high hydrogen enthalpy of 9.4e10.4 kJ/mol was estimated by a variant of the ClausiusClapeyron equation. These results underscored again the potential importance of such metal sites to hydrogen storage. Dailly et al.154 investigated the hydrogen sorption capacity of IRMOF-1 and IRMOF-8 at 77 K and high pressure. These two samples showed saturation adsorption capacity of 4.3 wt.% and 3.6 wt.%, respectively and the hydrogen sorption enthalpies for these two samples were calculated to be 6.1 and 4.1 kJ/mol. These values are consistent with previous results.135,143 Lin et al.155 synthesized and investigated the hydrogen storage properties for three copper co-ordination framework materials based on biphenyl, terphenyl, and quaterphenyl tetracarboxylic acids with formulas [Cu2((OOC)2C6H3$C6H3(COO)2)(H2O)2] (1), [Cu2((OOC)2C6H3$C6H4$ C6H3(COO)2(H2O)2] (2) and [Cu2((OOC)2C6H3$C6H4$C6H4C6H3(COO)2) (H2O)2] (3). They found that H2 adsorption was controlled by the available pore volume with a proportional decrease in adsorbate density with increasing pore size. The maximum gravimetric uptake of 7 wt.% was obtained for 3 with a surface area of 2932 m2/g resulting from its maximum pore volume of 1.284 cm3/g in that study (Figure 8.44). Whereas the maximum volumetric capacity of 43.6 kg/m3 was achieved for 2, with a surface area of 2247 m2/g because of a combination of substantial H2 uptake and the moderate density. Jhung et al.156 investigated the hydrogen storage properties of several nanoporous AlPOs, or aluminophosphates, to analyze the effect of pore diameter on hydrogen adsorption characteristics. They found that the heat of adsorption and adsorption capacity

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Figure 8.44 H2 sorption isotherms measured for desolvated 1 (circle), 2 (square) and 3 (triangle) at 78 K; the inset is an enlargement of the low-pressure region of the H2 isotherms; solid and open makers represent adsorption and desorption, respectively.155

per unit micropore volume increased with decreasing pore size. They suggested that the ideal porous material for efficient hydrogen storage was one with narrow pores and large micropore volume, or large surface area. Dietzel et al.157 synthesized, characterized, and investigated the hydrogen storage property of a novel three-dimensional coordination polymer Ni(dhtp)(H2O)2$8H2O [dhtp ¼ dihydroxyterephthe]. This novel sample con˚ , and possessed a Langmuir surface area of 1083 tained channels with a diameter of ~11 A 2 3 m /g and a pore volume of 0.41 cm /g. The maximum hydrogen adsorption for this sample was 1.8 wt.% at 77 K. The adsorbed hydrogen per Langmuir specific surface area corresponds to a value of 1.7 103 wt.% m2 g, which is similar to the value reported recently for the MOF Cu-BTC.143 The unsaturated metal sites within the coordination polymer can be attributed to this adsorption capacity. Frost et al.158 performed the GCMC simulations in a series of ten IRMOFs to predict the hydrogen adsorption isotherms. The results revealed the existence of three adsorption regimes: at low pressure (loading), hydrogen uptake correlated with the heat of adsorption; at intermediate pressure, uptake correlated with surface area; and at the highest pressure, uptake correlated with the free volume. The simulations also revealed that the 2015 DoE volumetric target of 81 kg/m3 might be difficult to meet using MOF materials, although there were a few materials that are either close to or even exceed the 2010 target of 45 Kg/m3. 122,150,155 Besides physisorption of hydrogen on the MOFs, hydrogen storage by spillover has been proposed as a mechanism to enhance the hydrogen storage density in MOF nanostructures. Li and Yang159 studied the hydrogen storage properties of a physical mixture of catalyst containing 5 wt.% Pt supported in active carbon, and IRMOF-1 or IRMOF-8 by employing the secondary hydrogen spillover technique. It was found that the spillover enhancement factors for IRMOF-1 and IRMOF-8 were 3.3 and 3.1; respectively, and the highest

Hydrogen Adsorption and Storage

measured reversible H2 uptake was 1.8 wt.% at 298 K and 10 MPa on a Pt/AC and IRMOF-8 physical mixture. In a further investigation160, they reported the hydrogen storage capacities of IRMOF-1 and IRMOF-8 by hydrogen spillover with carbon bridges (Figure 8.45). This bridge-building technique was effective for hydrogen uptake in MOF materials. The maximum hydrogen sorption capacity of 4 wt.% was obtained for the bridged IRMOF-8 at 298 K and 10 MPa. This was enhanced by a factor of 8 and 2 compared with the pure MOF materials and the physical mixture without bridges (Figure 8.46). In addition, the heat of adsorption for these two samples was measured to be in the range of 20e23 kJ/mol by using the Clausius-Clapeyron equation for the H2 adsorption isotherms measured at two different temperatures (298 and 323 K). From the above results, we conclude that MOFs are very promising materials for H2 storage at cryogenic temperature and moderate pressure (<100 bar). The building-block approach offers the capability of preparing frameworks with new topologies and finetuning both local and extended framework structures. A maximum adsorption of ~7.5 wt.% of hydrogen has been observed for MOF materials because of their high specific surface areas and small pore sizes. However, due to its relatively low density, the volumetric storage density of MOF is usually very low. To achieve the DoE 2015 goal, new strategies for the design of MOF nanostructures have high surface areas with pores of appropriate dimension, open metal sites, large free volumes, suitable crystal density, and large heats of adsorption are required.

Figure 8.45 (a) Primary spillover of atomic hydrogen from Pt metal to the activated carbon support and secondary spillover to the MOF receptor that has limited contacts with the support. (b) Facilitated primary and secondary spillover by using carbon bridges (black zone).160

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Figure 8.46 High-pressure hydrogen isotherms at 298 K for pure IRMOF-8 (,), Pt/AC and IRMOF-8 physical mixture (1:9 weight ratio;A) and for a bridged sample of Pt/AC-bridges-IRMOF-8: ﬁrst adsorption (B), desorption (:), and second adsorption (>).160

Hydrogen Storage of Zeolites Besides carbon nanostructures and MOFs, microporous zeolites are also believed to be promising candidates for hydrogen storage. They are of interest due to of the controllable diameter of the cages and channels, high stability, and low cost. One attractive feature of these microporous materials is that they contain isolated and exposed cations which could act as hydrogen binding sites. Weitkamp et al.161 reported hydrogen storage properties for a series of zeolites, with different pore structures and compositions at temperatures from 293 to 573 K and pressures from 2.5 to 10 MPa. Maximum adsorption capacity of 0.082 wt.% was demonstrated for zeolites with the highest concentration of sodalite cages at 573 K and 10 MPa. The authors implied that storage capacity was higher for the zeolites having a high number of small cavities in the structure. At cryogenic temperature, hydrogen uptake values have been reported to be much higher. Kazansky et al.162 investigated the hydrogen sorption properties of sodium forms of faujasite at 77 K, and found a hydrogen adsorption capacity of 1.2 wt.% at 77 K and 60 KPa. Takagi et al.163 measured hydrogen sorption properties of two zeolites, H-YZ and H-ZSM-5, together with several carboneous materials at 77 and 303 K over hydrogen pressures up to 3.5 MPa. A maximum hydrogen uptake of 0.6 wt.% was obtained for the H-YZ with a surface area of 710 m2/g at 77 K and 1 atm. It was also found that the amount of adsorbed hydrogen by weight depended on the micropore volume of the sample. Furthermore, Langmi et al.123 reported the hydrogen storage behaviors for zeolites A, X, Y, and RHO at temperatures ranging from 77 to 593 K and pressures up to 15 bar (Figure 8.47). A maximum hydrogen adsorption capacity of 1.81 wt.% was achieved for the zeolite NaY at 77 K and

Hydrogen Adsorption and Storage

Figure 8.47 Framework structures of zeolites: (a) zeolite A, (b) zeolites X and Y and (c) zeolite RHO. The corners on each framework represent Si or Al and these are linked by oxygen bridges represented by the lines on the frameworks.123

15 bar. The authors proposed that the hydrogen storage capacity correlated with the BET surface areas of zeolites in a near linear fashion (Figure 8.48). Later, the same group164 investigated the hydrogen storage properties of zeolites X, Y, A and Rho containing exchangeable cations. A maximum hydrogen sorption capacity of 2.19 wt.% was reported for CaX at 77 K and 1 atm. The authors claimed that the available void space of the frameworks was the major factor influencing the hydrogen uptake, which suggested that the cations may act as the binding sites for the hydrogen molecules. In another work, Vitillo et al.165 conducted theoretical studies for a total of twelve zeotypes by classic atomic simulations. Maximum amounts of hydrogen sorption ranging between 2.86 and 2.65 mass% were found. The predicted maximal storage capacities correlate well with experimental results obtained at low temperature. The

Figure 8.48 Correlation between hydrogen adsorption at 77K and BET surface area for a series of zeolite materials.123

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same group166 studied hydrogen adsorption for a series of Chabazite zeolite frameworks, H-SSZ-13, H-ZMS-5, and H-SAPO-34. The chabazite topology can be described as layers of double 6-rings that are interconnected by 4-rings. The double 6-ring layers stack in an ABC sequence. This leads to a framework with a regular array of barrelshaped cages interconnected by 8-ring windows. Figure 8.49 shows a schematic example for the dehydrated SSZ-13.167 The maximum hydrogen sorption capacity of 1.28 wt.% was found in H-SSZ-13 at 77 K and 0.92 bar. The adsorption energy of 10 kJ/mol, which is higher than that of activated carbon,58 was obtained for this framework. The authors suggested that the balance between hydrogen accessible space, high contact surface, and strong interaction with polarizing centers is important for hydrogen storage in this class of frameworks, and this viewpoint was further supported by another investigation.168 In 2005, Arean et al.169 studied the thermodynamic properties of hydrogen adsorption on the zoelites Na-ZSM-5 and K-ZMS-5 utilizing temperature infrared spectroscopy. It was found, consistent with the results of Zecchina and co-workers,166 that the adsorption enthalpies are 10 kJ/mol for H2/Na-ZSM-5 and 9.1 kJ/mol for the H2/K-ZSM-5 system. Recently, Li et al.170 investigated the hydrogen sorption properties of low silica type X zeolite LSX, which is fully exchanged by alkali-metal cations (Liþ, Naþ, and Kþ). The hydrogen adsorption capacity of LSX zeolite depended strongly on the cationic radius, and the density of the cations that are located on the exposed sites. Maximum hydrogen sorption capacities of 1.5 and 0.6 wt.% were obtained for a Li-LSX sample at 77 and 298 K, respectively. The hydrogen capacity in LSX zeolite

Figure 8.49 A ball and stick representation of the structure of dehydrated SSZ-13. The ﬁgure depicts the hydrogen atoms as spheres while the tetrahedral atoms (Al, Si) and oxygen atoms are omitted for clarity.167

Hydrogen Adsorption and Storage

by bridged hydrogen spillover was also investigated, and the hydrogen storage capacity was enhanced to 1.6 wt.% (by a factor of 2.6) at 298 K and 10 MPa. This represents the highest value for zeolite materials at room temperature although it is not as high as that of the AX-21101 and MOFs159,160 reported previously. The factor limiting the storage capacity of such systems is the relative large mass of zeolite frameworks which contain Si, Al, O, and heavy cations. Furthermore, many materials seem to have unnecessarily large void diameters.

Hydrogen Storage of Clathrates Trapping or encapsulation of hydrogen in microporous media is another strategy for hydrogen storage. The idea is to force hydrogen into the cavities of the nanostructured host. Upon cooling, hydrogen is trapped inside the cavities. It can be released again by raising the temperature. Thus, trapping may be a way to overcome the low binding energies of the physisorption process. The clathrate hydrogen hydrates were recently reported as potential hydrogen storage media through capturing the hydrogen in H2O cages.171 Clathrate hydrates are crystalline inclusion compounds, composed of a hydrogen-bonded water host lattice and guest molecules. There are three common types of gas hydrates structures: sI hydrate, which consists two 512 cages and six 51262 cages in a unit cell; sII hydrate, which consists of sixteen 512 cages and eight 51264 cages in a unit cell; sH hydrate, which consists of three 512, two 435663, and one 51268 cages in a unit cell. The crystal structures of these three types of clathrate hydrates are illustrated in Figure 8.50.172 The type of crystalline structure depends on the size of the guest molecule. Small guest molecules usually generate sI and sII hydrates and larger guest molecules resulted in the sH hydrate.171 Mao et al.173 reported a new type of sII clathrate hydrate with a molar ratio of H2 to H2O of approximately 1:2. This mixture was crystallized at a pressure of 220 MPa and 250 K. The clathrate cages formed were occupied with a cluster of two H2 molecules in the small cages and four in the large cages (Figure 8.51), resulting in hydrogen hydrate reversibly stored up to 5.3 wt.% hydrogen.174 The quenched clathrate is stable up to 145 K at ambient pressure. Raman spectroscopy showed the hydrogen molecules in the clathrate cages were still in free rotational states. However, a substantial softening and splitting of the vibron peaks of hydrogen indicated increased intermolecular interaction between hydrogen and water molecules in the cage.173 In another work, Patchkovskii and Tse175 performed a theoretical study of the thermodynamic stability of type sII hydrogen hydrate. They employed a statistical mechanical model in conjunction with first-principle quantum chemistry calculations. The theoretical analysis confirmed the experimental results, and suggested that the stability of the clathrate was primarily caused by dispersive interactions between H2 molecules and the water forming the cage walls. The calculation showed that the clathrate is thermodynamically stable at 2.5 MPa and 150 K. The main limitation of the

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(a)

(b)

(c)

Figure 8.50 Crystal structures of (a) sI hydrate; four unit cells viewed along a cubic crystallographic axis, (b) sII hydrate; two unit cells viewed along a face diagonal, and (c) sH hydrate; four unit cells viewed along the six-fold crystallographic axis (the positions of the hydrogen atoms have not been included). All cavities are assumed to be ﬁlled in these ﬁgures.172

Figure 8.51 (A) The sII crystal structure consisting of 51264 and 512 building blocks. (B) A tetrahedral cluster of four hydrogen molecules in the 51264 cage. (C) A cluster of two hydrogen molecules oriented toward opposite pentagonal faces in the 51264 cage. The H2O frameworks deﬁne the edges of the polyhedrons, with an oxygen atom at each corner and a hydrogen atom on each edge. The 51264 and 512 cages are drawn to scale relative to the enclathrated hydrogen molecule (sphere) of 2.72Å diameter.173

Hydrogen Adsorption and Storage

hydrogen clathrate is that the hydrogen cluster can only be stabilized in the cages at extremely high pressure. To store H2 at low pressure, Florusse et al.176 showed that the presence of THF (tetrahydrofuran), as a guest component, may indeed stabilize the clathrate at 280 K and 5 MPa. The PXRD (powder X-ray diffraction) measurements also confirmed that the same sII binary clathrate hydrate has formed as that reported by Mao and co-workers.173 Raman spectroscopy measurements suggested that all, or most of the large cages of the binary clathrate hydrate are filled with THF. In contrast, the H2 molecules most likely occupy to a large extent only the small cages. Although the inclusion of THF molecules can considerably decrease the synthetic pressure, their presence also leads to a significant decrease in the capacity for hydrogen storage. To solve this problem, Lee et al.177 reported that hydrogen storage capacities in THF-containing binary-clathrate hydrates can be increased up to 4 wt.% at modest pressures by tuning their composition to allow the hydrogen guests to enter both the larger and the smaller cages while retaining the low-pressure stability. The authors carried out a number of experiments at H2 pressure of 12 MPa and temperature of 277.3 K and found when the THF concentration was in the range of 2.0 to 5.56 mol%, H2 and THF molecules occupied the smaller and larger cages, respectively, leading to a storage capacity of 2.09 wt.% H2. However, a maximum storage capacity of about 4 wt.% H2 could be obtained when the concentration of THF was decreased to 0.15 mol%. In this case, H2 occupied not only all the small cages but also some of the larger ones (Figure 8.52). Further lowering the concentration of THF resulted in a failed formation of hydrogen hydrate under the conditions tested (12 MPa H2 and 277.3 K). Furthermore, Strobel et al.178 reported the hydrogen sorption properties of the binary THF-H2 clathrate hydrate. The maximum hydrogen storage capacity of this binary sII hydrate was found to be around 1 wt.% at moderate pressure (< 60 MPa). This result corresponds to the storage capacity for the stoichiometric binary hydrate with one hydrogen molecule per small cavity. Contrary to previous reports,177 hydrogen storage was not increased upon decreasing the THF concentration below the stoichiometric 5.6 mol% solution to 0.5 mol%, at constant pressure. These findings indicated that THF remains the favorable guest for the large cage, and that multiple hydrogen occupancies in this system cannot be achieved at moderate pressures. More recently, Chapoy et al.179 investigated hydrogen storage in semi-clathrate hydrates of quaternary ammonium compounds. The synthesized binary H2-TBAB (TBAB ¼ (C4H9)4NþBr) and H2TBAF (TBAF ¼ (C4H9)4NþF) semi-hydrates are stable at ambient temperature and pressure. In addition to favorable stability, the hydrogen adsorption capacities of these hydrates are an order of magnitude greater than those for previously reported binary H2-THF clathrates178 at the same pressure. Obviously, this system is strongly sensitive to promoters, leaving room for further development by optimization of the promoter system. But the slow kinetics of the current synthesis approach for the hydrogen clathrate also creates difficulties in terms of

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Figure 8.52 Hydrogen gas content as a function of THF concentration and a schematic diagram of H2 distribution in the cages of THF þ H2 hydrate. In region III, H2 molecules are stored only in small cages, while in region II both small and large cages can store H2 molecules. At the highly dilute THF concentrations of region I, H2 molecules can still be stored in both cages, but extreme pressures (2000 bar) are required to form the hydrates.177

practical application. Thus, further investigation is needed to develop a hydrogen clathrate system with high storage capacity and fast kinetics properties.

CHEMISORPTION OF HYDROGEN Chemisorption is a type of adsorption whereby a molecule adheres to a surface through the formation of a chemical bond. This process is characterized by high temperatures, strong interaction, and high enthalpies. Chemisorption materials may provide high volumetric and gravimetric storage capacities. However, due to the strong interaction between the adsorbate and surface, reversibility can only be observed at high temperature.

Hydrogen Storage of Metal Hydrides Metallic hydrides form a wide range of stoichiometric and non-stoichiometric compounds by direct interaction of hydrogen with metals. Charging can be

Hydrogen Adsorption and Storage

accomplished using molecular hydrogen gas, or hydrogen atoms from an electrolyte. The detail process is illustrated schematically in Figure 8.53.1 In the case of gas phase loading, molecular hydrogen is adsorbed on the surface, and then dissociated at the surface before absorption. While, in the desorption process, two H atoms recombine to molecular hydrogen (R-8.1). x M þ H2 5MHx 2

(R-8.1)

where M ¼ metal or an intermetallic compound MHx ¼ hydride x ¼ ratio of hydrogen to metal (H/M). When hydrogen molecules approach a metal surface, a physisorbed film is first formed on the metal surface by Van der Waals forces. In this process, a hydrogen gas molecule interacts with several atoms at the surface of the metal. The physisorption energy is typically of the order 10 kJ/mol. The potential energy of the molecule shows a minimum at approximately one molecule radius. The interaction consists of an attractive term, which diminishes with the distance between the hydrogen molecule and the surface of the metal to the power of 6, and a repulsive term diminishing with the distance to the power of 12. Therefore, in the next step of hydrogenemetal interaction, the hydrogen molecule has to overcome an activation barrier for dissociation (H2/2H, ED ¼ 435.99 kJ/mol), and to form a chemisorbed state at the surface of the metal before diffusion into the bulk metal (see Figure 8.54).2 The chemisorption energy is typically of the order of 50 kJ/mol. After dissociation on the metal surface, the hydrogen atoms generally diffuse

Figure 8.53 Schematic model of a metal structure with H atoms in the interstices between the metal atoms, and H2 molecules at the surface. Hydrogen atoms are from physisorbed hydrogen molecules on the left-hand side and from the dissociation of water molecules on the right-hand side.1

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Figure 8.54 Schematic of potential energy curves of hydrogen approaching a metal.2

rapidly through the bulk metal, even at room temperature, to form an M-H solid solution, commonly referred to as a-phase. The thermodynamic aspects of H2 hydride formation from gaseous hydrogen are described by pressureecomposition isotherms (Figure 8.55).2 The hydrogen concentration in the a-phase (CH) is pressure dependent and frequently follows Sieverts’ law: CH ¼ KP 1=2

Eqn. 8.22

where K is a temperature dependent constant and P is the dihydrogen pressure. As the hydrogen pressure does not increase together with the concentration of H in the metal, locally interactions between hydrogen atoms become important. Nucleation and growth

Figure 8.55 Typical isothermal pressure-composition (PCT) curves for an ideal metal-hydrogen system.2

Hydrogen Adsorption and Storage

of the hydride (b) phase will begin. While the two phases coexist, according to the Gibbs phase rule, the conversion of the saturated solid solution to hydride should take place at a constant pressure. This will show a flat plateau in a PCT curve. The temperaturedependent pressure is the equilibrium dissociation pressure of the hydride. And the length of the flat plateau determines how much can be stored reversibly with small H2 pressure variations. In the pure b-phase, the pressure rises steeply with the H2 concentration. At higher pressure, additional plateau and hydride phases may be formed. The two-phase region ends in a critical point TC, above which the transition from a- to b-phase is continuous. The plateau or equilibrium pressure depends strongly on temperature, and is related to the changes of enthalpy and entropy, respectively. The relationship can be given by the van ’t Hoff equation: Peq DH 1 DS lnð Þ ¼ : R T R Peq

Eqn. 8.23

where DS ¼ entropy change DH ¼ enthalpy change R ¼ gas constant T ¼ temperature. Because the entropy change corresponds mostly to the change from molecular hydrogen gas to dissolved solid hydrogen, it is approximately the standard entropy of hydrogen (S0 ¼ 130 J$K1 mol1) and is, therefore, DS z -130 J$K1 mol1 H2 for all metal-hydrogen systems. The enthalpy term characterizes the stability of the metalehydrogen bond. To reach an equilibrium pressure of 1 bar at 300 K, DH should be 39.2 kJ mol1 H2. The entropy of formation term for metal hydrides leads to a significant heat evolution DQ ¼ T$DS (exothermal reaction) during hydrogen absorption and the same heat has to be provided to the metal hydride to desorb the hydrogen (endothermic reaction). If the hydrogen desorbs below RT, this heat can be delivered by the environment. However, if the desorption is carried out above RT, the necessary heat has to be delivered from an external source, such as the combustion of hydrogen.2 However, most practical hydriding metals do not show a perfectly flat plateau and zero hysteresis. The most common isothermal PCT curve is a hysteresis loop as shown in Figure 8.56; which clearly shows the mathematical and numerical definitions of hysteresis, plateau slope and different hydrogen adsorption capacities.180 It has been reported that approximately fifty among ninety-one metallic elements in the periodic table (I and II main groups and III, IV and V transition groups of metals) alloys are able to absorb hydrogen (see Table 8.2).181 However, most elements are not conveniently used at 1e10 atm, 0e100 C range for practical applications, as shown by the van ’t Hoff lines of Figure 8.57.180

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Figure 8.56 Schematic isothermal pressureecomposition hysteresis loop.180

As shown in Figure 8.57, only vanadium is in this range and solid solutions of V and other metals are of current interest. Pd has been used for more than 100 years for H-storage, but it is very expensive, does not hold much H and requires heating well above 100 C to liberate that H. None of the elements is currently suitable for practical Table 8.2 Compilation of known hydrides for the transition metals and groups IA, IIA and IIB. Li

BE

LiH (rocksalt)

BeH2

NA

Mg

NaH (rocksalt)

MgH2 (rutile) MgH2 (fluorite) at high P

K

Ca

Sc

KH (rocksalt)

CaH2 (Co2Si)

ScH2 (TiH2 V2H, V2H3 (fluorite) (fluorite) VH (H interslitiel TiH2 solution in bcc V) (ThH2) at low T VH2 (fluorite)

Ti

CrH (NiAs) MnH at high P (NiAs)

Mo

Zr

V

Nb

Cr

Mn

CrH2 (fluorite)

Fe

Co

Ni

Cu

Zn

FeH CoH, CuH NiHo2 CoH2 (wurtzite) (double (nonhexagonal) P> 50 MPa stoichiometric at high P, rocksalt) low T

Unstable hydrides

Ru

Cd

CrH3

Rb

Sr

Y

RbH (rocksalt)

SrH2 (C02Si)

YH2 ZrH2 Nb2H (H MoH (fluorite) (fluorite) interslitiel (NiAs) at in Nb) high P YH3 NbH ZrH2 (struc. (BiF3) (ThH2) at low T unknown) NbH2 (fluorite)

Tc

Rh

Pd

RhH (CsCI) at high P

Pd2H2 PdH4 (NiMo4)

Ag

Unstable hydrides

Hydrogen Adsorption and Storage

Figure 8.57 Van ’t Hoff lines (desorption) for elemental hydrides. Box indicates 1e10 atm, 0e100 C ranges.180

applications.180 Since metal hydrides containing only one metal have limited practical applications in hydrogen storage due to of their high thermodynamic stabilities; the metallic hydrides of intermetallic compounds have been investigated over the past several decades in order to find a material that meets the practical requirements. Some of the metallic hydrides of interest for hydrogen storage purposes are listed in Table 8.3.182 The prototype intermetallic hydrides are composed of two elements. The A element is usually a rare earth or an alkaline earth metal, and tends to form a stable hydride. The B element is often a transition metal and forms only unstable hydrides, such as Pt and Ru, which can absorb considerable quantities of hydrogen; together with Pd and Ni, which are excellent hydrogenation catalysts although they do not form hydrides.2 Table 8.3 The most important families of hydride-forming intermetallic compounds. Intermetallic compound Prototype Hydrides Structure

AB5 AB2 AB3 A2B7 A6B23 AB A 2B

LaNi5 ZrV2, ZrMn2, TiMn2 CeNi3, YFe3 Y2Ni7, Th2Fe7 Y6Fe23 TiFe Mg2Ni, Ti2Ni

LaNi5H6 ZrV2H5.5 CeNi4H4 Y2Ni7H3 Ho6Fe23H12 TiFeH2 Mg2NiH4

Haucke phases, hexagonal Laves phase, hexagonal or cubic Hexagonal, PuNi3-type Hexagonal, Ce2Ni7-type Cubic, Th6Mn23-type Cubic, CsCl- or Ti2Ni-type Cubic, MoSi2- or Ti2Ni-type

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Typical alloys for hydrogen storage are the AB5 compounds such as LaNi5 with a hydrogen content of roughly 1.4 wt.%, and an equilibrium pressure around 2 bar at room temperature.183 However, its cost is relatively high. The P-C-T diagram shows a flat plateau and low hysteresis, but unfortunately the hydrogen capacity is degraded after a few cycles. Therefore, these materials are unattractive for on-board hydrogen storage. Another interesting type of intermetallic alloy for hydrogen storage is AB2, which represents a large and versatile group of hydriding materials (nearly 500 types). The A-elements are often from the IVA group (Ti, Zr, Hf) and/ or rare earth series (at. no. 57e71) or Th. The B-elements can be a variety of transition or non-transition metals with something of a preference for atomic numbers 23e26 (V, Cr, Mn, Fe). A very wide variety of substitutions are possible for both A- and B-elements, thus providing a high degree of fine tuning of PCT properties.180 However, only approximately 2 wt.% of hydrogen storage capacity can be obtained for this type of compound.184 This is still far short of meeting the US DoE goal of 9 wt.% reversible hydrogen capacity. These materials are cheaper compared to the LaNi5 alloys but more sensitive to gaseous impurities. Thus, a small amount of oxygen can be a poison for the AB2s, while AB5s it acts as a reactant, reducing the storage capacity slightly. Compared with above two types of intermetallic alloys, A2B-type compounds have demonstrated higher hydrogen storage capacity, e.g. 3.6 wt.% for Mg2NiH4. However, the high desorption temperature (close to 300 C) of Mg2NiH4185e188 is too high for most on-board applications. Metallic hydrides usually have extremely high volumetric density. For example, a volumetric density of 115 kg/m3 was reached in LaNi5H6.182 However, all the reversible hydrides working around ambient temperature and pressure consist of the heavy transition metals; therefore, the gravimetric density of hydrogen is usually limited to less than 3 wt.%. For example, LaNi5H6 has a gravimetric hydrogen density of 1.4 wt.%. Therefore, light metal hydrides such as Mg and Ca appear to be promising candidates for hydrogen storage. In fact, MgH2 containing 7.6 mass% hydrogen has already been reported. It offers the benefit of low cost and abundant availability of Mg and good reversibility. But its formation from bulk Mg and gaseous hydrogen is extremely slow. Other disadvantages of MgH2 as a hydrogen storage material are the slow desorption kinetics, high temperature of hydrogen discharge, and a high reactivity towards air and oxygen. To increase the sorption kinetics, nanocrystalline Mg hydrides can be prepared by a mechanochemical process such as high-energy ball-millling.189,190 Once ball-milled, the adsorption and desorption kinetics can be enhanced,191,192 as shown at Figure 8.58.192 Figure 8.58 shows that the desorption of unmilled MgH2 is very slow. There is no obvious desorption at 573 K within 2000 s. Even at 623 K, it takes more than 3000 s to desorb completely. In comparison, milled MgH2 can fully desorb at 623 K in about 700 s. Although the desorption at 573 K is slower, it can still desorb about 2.2 wt.% of hydrogen within 2000 s. The nucleation and growth process of MgH2 takes place during

Hydrogen Adsorption and Storage

Figure 8.58 Hydrogen desorption curves of unmilled MgH2 (solid symbols) ball-milled (open symbols) MgH2 under a hydrogen pressure of 0.015.192

desorption. This is shown by the sigmoidal desorption curve in Figure 8.58. This is the case in both unmilled and ball-milled MgH2. By fitting the desorption curves with the Johnson-Mehl-Avrami (JMA) equation: a ¼ 1 expfðKtÞn g

Eqn. 8.24

where n ¼ reaction exponent a ¼ desorption fraction at time t K ¼ K0 exp(Q/RT) Q ¼ activation energy R ¼ gas constant T ¼ temperature. The reaction exponent (n) in this study was found to be 3 for both unmilled and milled powder. This value can be interpreted as an instantaneous nucleation followed by an interface-controlled three-dimensional growth process. Using the JMA equation (Eqn. 8.24), the rate constant K can also be determined by fitting the desorption curve. From the Arrhenius plot of K values with temperature, the activation energies of desorption (R) for the ball-milled and unmilled MgH2 are found to be 120 and 156 kJ/mol H2, respectively.192 To further improve the H exchange kinetics, composite materials were prepared by ball milling of Mg with a series of transition metals e.g. V,193e195 Nb,196

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Pd,197 Pt,198 Ni,188,197e200 Ru,198 Fe,195,200e203 Ti,195 Mn,195 Co.200,204 The fast kinetics were explained by a variety of factors such as: reduction of activation energy, introduction of defects by intensive milling, increase of specific surface area, small effective particle size, and particular micro-structure of the composite. The main disadvantages of metal hydrides are their low gravimetric densities for those that can operate at room temperature; whereas for some metallic hydride systems with a relatively high gravimetric storage capacity such as MgH2, high desorption temperature, and slow hydriding and dehydriding rates are the main challenges for practical storage applications. Thus, new catalysts and proper preparation methods are needed to overcome these problems and obtain practical materials.

Hydrogen Storage of Complex Hydrides Group I and II salts of [AlH4] and [BH4] (alanates and borohydrides) have recently received considerable attention as potential hydrogen storage materials.205 All of these materials are currently referred to as “complex hydrides”. Actually, all the hydrides with hydrogen to metal ratio of more than two are ionic or covalent compounds, and belong to this complex hydride group. The main difference between complex and metallic hydrides is the transition to an ionic, or covalent compound upon hydrogen absorption. The hydrogen in the complex hydrides is often located in the corners of a tetrahedron with B or Al in the center. The negative charge of the anion, [BH4] and [AlH4], is compensated for by a cation, typically an alkali (e.g. Li), alkaline earth metal (e.g. Mg), or a transition metal (e.g. Zn). The storage capacity thus depends on the weight of the metals; and the number of possible bound hydrogen atoms in conjunction with the charge of the complex anion. Ideally, light metals with high vacancy are desired for high storage capacity. Most complex hydrides are known to be stable and decompose only at elevated temperatures, often above the melting point of the complex.182 However, the decomposition temperature can be reduced, and the hydrogen adsorption efficiency can be improved by using various dopants or reducing the grain size. Some of the recent developments in alanates and borates are reviewed in the next section. Alanates The lightweight complex aluminum hydrides have received the most attention over the past decade, especially NaAlH4 due to its high reversible hydrogen storage capacity at moderate temperatures. Bogdanovic and Schwickardi206 showed, for the first time, the reversible hydrogen storage system of NaAlH4 doped with a small amount of Ti compound. A maximum hydrogen storage capacity of 4.2 wt.% was obtained at 210 C. Despite the slow kinetics, the discovery of reversibility in the complex hydrides was a breakthrough in solid-state hydrogen storage. In a further study,207 they found that through variation of NaAlH4 particle sizes, catalysts (dopants), and doping procedures; kinetics as well as the de- and rehydrogenation stabilities within different cycles can be

Hydrogen Adsorption and Storage

substantially improved. Thermal decomposition of NaAlH4 at higher temperatures took place in two steps to give NaH, Al, and H2. In principle the first step can give 3.7 wt.% H2 and up to 5.5 wt.% can be achieved for the total two steps: 3 NaAlH4 /Na3 AlH6 þ 2 Al þ 3 H2 ð3:7 wt:% HÞ

(R-8.2)

Na3 AlH6 / 3 NaH þ Al þ 3=2 H2 ð1:8 wt:% HÞ

(R-8.3)

The first decomposition step has an equilibrium pressure of 0.1 MPa at around 30 C; which means that this material is comparable to a typical low-temperature metal hydride. The hydrogen content for this first decomposition step (3.7 wt.% H2) is higher than a classic low-temperature metal hydride by a factor of two. The second decomposition step releases 1.8 wt.% of hydrogen. The equilibrium pressure of this step is 0.1 MPa at about 100 C, which is typical for a medium-temperature metal hydride. Furthermore, the enthalpies for the dissociation reactions were determined to be 37 and 47 kJ/mol for the first and second dissociation steps of Ti-doped NaAlH4, respectively. However, one disadvantage of Na alanate was its ineffective rehydrogenation, which required very harsh conditions. More recently, Bogdanovic et al.208 reported the hydrogen storage properties of Na alanate doped with titanium nanoparticles and found when using TiN as a doping agent 4.9e5.1 wt.%, which is close to the theoretical limit of 5.5 wt.%, was obtained at 104 or 170 C in the pressure range of 140e115 bar. In addition, for the colloidal Ti particle doped NaAlH4 hydrogenation times of 10e15 minutes, approaching those required for practical applications, combined with high capacities of 4.5 wt.% have been realized in cyclical tests (Figure 8.59).

Figure 8.59 Hydrogen storage capacity in a 25 cycles test of NaAlH4 doped with 2 mol% of colloidal Ti particles. Dehydration at 120/180 C and normal pressure; hydrogenation at 100 C and 100e85 bar.208

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As shown at Figure 8.59, after cycles 2e5, a storage capacity of 4.8 wt.% H2 was exhibited. In the following cycles the capacity remained at between 4.5 and 4.6 wt.% H2 throughout of the test. On the other hand, the hydrogenation rate showed a decreasing trend. Thus, further research efforts aimed at improving the stability of the hydrogenation rate and of hydrogenation under low pressure conditions are needed. Another frequently used doping precursor is TiCl3. The active catalytic species is produced during the treatment of the material in a ball-milling procedure. During ball milling, the dopant is highly dispersed within the hydride material. This is necessary for good kinetics in solid-state reactions. Sandrock et al.209 studied the reversible hydrogen storage properties of Ti-catalyzed NaAlH4 as a function of Ti content using a dry preparation technique consisting of the ball-milling of NaAlH4 þTiCl3 mixtures (0e9 mol% TiCl3). The general solid state reaction was proposed as: ð1 xÞNaAlH4 þ xTiCl3 /ð1 4xÞNaAlH4 þ 3xNaCl þ xTi þ 3xAl þ 6xH2 (R-8.4) where x is the mole fraction of TiCl3 added to the initial ball milling charge. This reaction can be further simplified as: 3NaAlH4 þ TiCl3 /Ti þ 3Al þ 3NaCl þ 6H2

(R-8.5)

According to R-8.4 and R-8.5, the TiCl3 was reduced to metallic Ti; which serves as a powerful catalyst for enhancing both dehydrating and hydriding kinetics even at the lowest doping level (0.9 mol%). They also found that as Ti content increases, the hydriding and dehydrating kinetics increased. Although Ti-catalyzed alanates have reasonably good charging rates at 125 C and discharge from NaAlH4 to Na3AlH6 is also fast enough at the same temperature, the decomposition for the Na3AlH6 to NaH is relatively slow. Thus, improvements of the catalyst are still needed for practical charging and discharging kinetics below 100 C. In another work, hydrogen absorption and desorption of NaAlH4 doped with TiCl3 were studied by dynamic in situ X-ray diffraction under conditions similar to those found in fuel cell operations.210 Catalyst doping was found to dramatically improve kinetics under these conditions. XRD measurements showed that TiCl3 reacted with NaAlH4 to form NaCl, and most likely zero valence Ti. Furthermore, the addition of TiCl3 promoted the partial decomposition of NaAlH4 to Na3AlH6 during the doping process. More recently, X-ray absorption spectroscopy was applied to investigate NaAlH4 doped with 5 mol% TiCl3 by ball milling.211 XANES (X-ray absorption near-edge structure) analysis confirmed that after the ball milling Ti species were reduced from Ti (þ3) to Ti (0). Release or absorption of hydrogen did not affect the chemical state obtained after ball milling, which is in good agreement with the results reported by Felderhoff and co-workers.212 EXAFS (extended X-ray absorption fine structure) analysis showed that the Ti atoms were associated only with Ti as next neighbors in the ball-milled state, as well as during subsequent desorption

Hydrogen Adsorption and Storage

and absorption of hydrogen. Furthermore, an increase of the particle size and an ordering of the local structure were seen to evolve with the desorption and the absorption of hydrogen. Zidan et al.213 investigated the dehydrogenation kinetics and hydrogen cycling behaviors of Ti and Zr doped NaAlH4. They found that the dehydriding kinetics of NaAlH4 was significantly enhanced through zirconium doping , but Zr was inferior to Ti as a catalyst for the dehydriding of NaAlH4 to Na3AlH6 and Al (see R-8.2), but a superior catalyst for the dehydriding of Na3AlH6 to NaH and Al (see R-8.3). Thus, titanium and zirconium can act in concert to optimize the dehydriding/ rehydriding behavior of NaAlH4. After the initial dehydriding/rehydriding cycle, NaAlH4 which was doped with titanium and/or zirconium, was stabilized with a greater than 4 wt.% cyclable hydrogen capacity. The onset of rapid dehydriding occurred in the titanium-containing materials at temperatures below 100 C. Thomas et al.214 observed significant changes in the particle morphology and elemental distribution during the hydrogen desorption and cycling process of Ti-Zr catalyzed NaAlH4 by employing SEM (scanning electron microscopy) and EDS (energy dispersive spectroscopy). More important, their results indicated that the initial dehydriding reactions were accompanied by significant enhancement of Al concentration toward the surface of particles and that elemental segregation of Na and Al, which occurred with repeated absorption/ desorption cycles. This could contribute to the decreased hydrogen storage capacity over the course of several cycles. Furthermore, Anton215 reported the effect of a wide range of transition metal and rare earth element dopants with various valance states on the hydrogen sorption capacity and hydrogen discharge kinetics of NaAlH4. The amount and type of dopant had a substantial effect on these parameters. However, only metal cations, not the anions, of the catalysis played a role in the dehydrogenation kinetics and the most active cation additions were found to be Ti4þ. Moreover, Jensen et al.216 reported the hydrogen sorption properties of catalyzed NaAlH4, in which Ti was introduced by a novel dry doping method under an air- and water-free environment. It was found that the dehydrogenation of this material occurred about 30 C lower than that found previously (Figure 8.60).206 This kinetic enhancement was undiminished over several dehydriding/hydriding cycles. The results also indicated that only a fraction of the titanium introduced into the materials was catalytically active. Balde et al.217 further investigated the catalytic effect of Ti-species in TiCl3-doped NaAlH4 by EXAFS, XANES and XRD. They found that in NaAlH4 ball-milled with TiCl3, the majority (~70%) of Ti occupied interstitial spaces in the NaAlH4 lattice. The remaining Ti was present at the surface of Al. After desorption at 125 C the Ti present at the Al surface appeared to be the most active Ti species for hydriding catalysis of a desorbed NaAlH4. At 225 C, Ti migrated from the Al surface to the bulk, forming amorphous TiAl3 clusters. This surface to bulk migration of the Ti atoms was accompanied by a deactivation of the catalyst. Subsequently, the TiAl3 clusters agglomerated to crystalline TiAl3 during heat treatment to 475 C, which lead to a lower dispersion of the

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Figure 8.60 Thermal programmed desorption (2 C min1) of hydrogen from undoped, wet titanium doped, and dry titanium doped NaAlH4.216

Ti catalyst and consequently a lower hydriding activity. Thus, the rank of the Ti activity for hydrogen absorption was: Ti on the Al surface > TiAl3 cluster > crystalline TiAl3. Zaluska et al.218 investigated the hydrogen sorption properties of both NaAlH4 and Na3AlH6 by mechanical grinding and chemical modification. They found that milled NaAlH4 or Na3AlH6 exhibited greater enhancement of the kinetics of adsorption and desorption. The addition of carbon in the milling process improved their performance even more remarkably. Mixtures of NaAlH4 with carbon were capable of reversible dehydrogenation/ hydrogenation at much lower pressures and with much faster kinetics than conventional compounds (Figure 8.61). Not only can the Ti doping accelerate the hydrogenation and dehydrogenation reactions in the Ti-NaAlH4 system, but it also significantly alters the thermodynamics of the system. This is demonstrated by the change of the dissociation pressure with doping level (Figure 8.62).219 During the dehydrogenation reaction, the highly concentrated TiAl3 alloy produced in the doping reaction was diluted by the aluminum, which was generated during the reaction. It was this dilution process that gave an additional contribution to the free energy of the system and destabilized it at higher dopant concentrations. In addition to NaAlH4, magnesium alanate (Mg(AlH4)2) is another interesting complex hydride. It exhibits high gravimetric hydrogen density although it is less stable than sodium alanate. Fichtner et al.220 synthesized Mg(AlH4)2 in a metathesis reaction of magnesium chloride and sodium alanate in ether solution according to R-8.6 MgCl2 þ 2NaAlH4 /MgðAlH4 Þ2 þ 2NaCl

(R-8.6)

Thermal analysis of synthesized magnesium alanate showed decomposition with a release of hydrogen proceeding in two major steps (R-8.7 and R-8.8).

Hydrogen Adsorption and Storage

Figure 8.61 Decomposition of NaAlH4 at 160 C: (a) ball-milled for 2 h without carbon, ﬁrst decomposition; (b) ball-milled for 2 h with 10 wt.% of carbon, ﬁrst decomposition; (c) ball-milled with 10 wt.% of carbon, second decomposition.218

MgðAlH4 Þ2 /MgH2 þ 2Al þ 3H2

(R-8.7)

MgH2 /Mg þ H2

(R-8.8)

The peak decomposition temperature of the first step was found to be 163 C. The residue at 200 C consisted of MgH2 and Al, which continued to release hydrogen and transformed into an Al3 Mg2/Al mixture at higher temperatures, according to R-8.9. 1 1 2Al þ Mg/ Al3 Mg2 þ Al 2 2

(R-8.9)

In the first decomposition step, 6.6 wt.% of hydrogen was released compared to the theoretical total hydrogen storage capacity is 9.3 wt.%. The main difference in the decomposition of Mg (AlH4)2 compared to Na alanates, is that it transforms into a nonalanate material during the first decomposition step. This has been confirmed by an in situ synchrotron X-ray diffraction study.221 By adding a TiCl3 promoter and ball milling the samples, the peak temperature could be reduced by up to 45 C. However, this was accompanied by a reduced hydrogen release. The authors indicated that even though Mg (AlH4)2 had promising high hydrogen storage capacity, some important issues needed to be explored before practical application. These issues included the slow thermodynamic hydrogen absorption and release and the formation of stable MgH2 and Al2Mg2, which negatively influenced the re-absorption of hydrogen. Komiya et al.222 further investigated the hydrogen dehydriding and rehydrogenation behavior of TiCl3 doped Mg (AlH4)2. They found that this system was decomposed in the two-step reactions, which

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Figure 8.62 PCI curves for 0.5, 2, 4, 10, 17.5 and 25 mol% Ti-doped NaAlH4 at 160 C.219

is in agreement with Fichtner et al.’s result.220 The decomposition temperature of the first reaction could be reduced remarkably when it was mixed with TiCl3. The decomposition temperature decreased down to ambient temperature while increasing content of TiCl3 in the specimen. However, the re-hydrogenation reaction barely occurred even in the pre-decomposed TiCl3-containing specimen. For example, only 0.3 mass% of hydrogen was absorbed in the 1 mol% TiCl3-doped specimen by exposing it for a long period to the hydrogen atmosphere of 30MPa at 353 K. This result indicated that Mg (AlH4)2 was too unstable to result in smooth re-hydrogenation. This has been demonstrated by a recent DSC study of decomposition of Mg(AlH4)2223, in which the dissociation of Mg(AlH4)2 to MgH2, Al and hydrogen was accompanied by a heat uptake of only ~1.7 kJ/mol. The thermodynamic stability of Mg(AlH4)2 is, therefore, far below that suitable for the reversible hydrogen storage (15e24 kJ mol H1).1 Furthermore, in another study, Kim et al.224 synthesized Mg(AlH4)2 by a mechanochemically activated metathesis reaction between NaAlH4 and MgCl2 without solvent. Since the assynthesized Mg(AlH4)2 does not contain any solvent adduct, this strongly influences the thermal decomposition behavior of magnesium alanate, its thermal decomposition can be described clearly. Through TG/MS (thermogravimetric analysis and mass spectrometry) and DSC (differential scanning calorimetry), the thermal decomposition behavior of solvent-free Mg(AlH4)2 was analyzed. It has been confirmed that Mg(AlH4)2 decomposes in two steps as previously reported220 (R-8.7 and R-8.8). The first

Hydrogen Adsorption and Storage

Figure 8.63 TPD (5 C/min) of the (A) NaAlH4, (B) LiAlH4, and (C) Mg(AlH4)2 systems when doped with Ti and ball milled.225

decomposition reaction is exothermic and starts at about 115 C (R-8.7). The second decomposition is endothermic in total and begins at about 240 C (R-8.8). Due to the fact that active Al formed after the first decomposition step destabilized MgH2 to produce more stable Al3Mg2 (R-8.9), the authors draw the conclusion that Mg(AlH4)2 was not likely to be irreversible under practical conditions. Wang et al.225 provided the first comprehensive comparison of the effect of Ti as a dopant on the dehydrogenation of NaAlH4, LiAlH4 and Mg(AlH4)2 complex hydrides. Figure 8.63 displays the typical behavior of the dehydrogenation of NaAlH4, LiAlH4 and Mg(AlH4)2 doped with different amounts of Ti during TPD. Figure 8.63A shows that NaAlH4 has a higher hydrogen release rate and lower temperature with increased Ti doping level (3 wt.% hydrogen can be released at 140 C). Figure 8.63B ilustrated, the effect of the Ti dopant is very pronounced. Increasing the dopant level causes hydrogen to be released at a much lower temperature, but also in smaller amounts. The LiAlH4 system, which is very attractive, can release 3 wt.% hydrogen at 100 C. However, this system is irreversible at

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125 C and 1,200 psig. Figure 8.63C shows the pronounced effect of Ti on the hydrogen release rate, but also indicates, in agreement with Fichtner et al.’s result,220 that increased Ti doping level decrease the hydrogen release capacity. The Mg(AlH4)2 system can release 6 wt.% hydrogen at 150 C; however, this dehydrogenation temperature is still too high for most practical applications and reversibility is still questionable. Borohydrides Recently, borohydrides have attracted significant attention for hydrogen storage due to their high hydrogen content. For example, the overall hydrogen amount of LiBH4 is 18.5 wt.%, much higher than in the alanate system. In general, the alkali borohydrides have a large decomposition enthalpy (too stable), and require rather high temperature for hydrogen release. Current research efforts are focused on reducing the decomposition enthalpy and improve the reaction kinetics of borohydride systems. Schlesinger et al.226 first reported the synthesis of a pure LiBH4 by the reaction of ethyllithium with diborane (B2H6) (R-8.10). 3LiC2 H5 þ 2B2 H2 /3LiBH4 þ ðC2 H5 Þ3 B

(R-8.10)

The direct reaction of the Li hydride with diborane in etheral solvents under suitable conditions produces high yields of the LiBH4 (R-8.11).227 2LiH þ B2 H6 /2LiBH4

(R-8.11)

Direct synthesis from the metal, boron, and hydrogen at 550e700 C and 3e15 MPa H2 has been reported to yield the lithium salt (R-8.12). It is claimed that such a method is generally applicable to groups IA and IIA metals.228 M þ B þ 2H2 /MBH4

(R-8.12)

where M ¼ Li, Na, K, etc. Fedneva et al.229 investigated LiBH4 by DTA (differential thermal analysis). The thermogram of LiBH4 showed three endothermic effects: at 381e385, 541e559, and 756e765 K. The endothermic effect at 381e385 K was reversible and corresponded to a structural transition from orthorhombic to polycrystalline. The second peak at 541e559 K corresponded to LiBH4 fusion which is accompanied by a slight decomposition. This liberates approximately 2 wt.% of the hydrogen from the compound. The main evolution of gas started at 653 K and released 80% of the hydrogen from LiBH4. The cause of the small effect at 756e765 K was uncertain, but it coincided with the liberation of 50% of the theoretical hydrogen content in LiBH4. Muller et al.230 reported that LiBH4 decomposed and released 13.8 wt.% hydrogen at 723 K and 1.3 Pa within 24 hours by the reaction: 3 LiBH4 /LiH þ B þ H2 (R-8.13) 2

Hydrogen Adsorption and Storage

Figure 8.64 Generalized enthalpy diagram illustrating destabilization through alloy formation upon dehydrogenation. Including the alloying additive, B, reduces the enthalpy for dehydrogenation through the formation of ABx and effectively destabilizes the hydride AH2.231

The hydrogen desorption enthalpy was found to be 24.8 kcal/mol. After dehydrogenation, the sample was able to absorb 11.8 wt.% of hydrogen at 923 K and 15 MPa within 48 hours. To improve the dehydrogenation and subsequent rehydrogenation processes, 10 wt.% aluminum powders were mixed mechanically with LiBH4. The resulting material then released 12.4 wt.% hydrogen at 723 K and 1 Pa within 24 hours. Then, 15.2 wt.% hydrogen was recharged onto the decomposed material at 923 K and 15 MPa within 12 hours. However, this extremely slow reversibility lasted for only two cycles. The dehydriding and rehydriding temperatures were too high. To reduce decomposition enthalpy of borohydrides, an approach that involves incorporating a second species into the reaction to stabilize the reaction product has been proposed (Figure 8.64).231 From Figure 8.64, in isolation, the pure hydride AH2 undergoes dehydrogenation to form AþH2 with a relatively high enthalpy. Consequently, the equilibrium hydrogen pressure will be low. Alternatively, the temperature required for the equilibrium pressure (T (1 bar) in Figure 8.64) will be high. However, if the chemical environment of AH2 is altered by adding a second component, B, which alloys with A, then dehydrogenation can proceed to ABxþH2. This reaction occurs with a reduced enthalpy and, therefore, an increased equilibrium hydrogen pressure and lower desorption temperature. Thus, AH2 is effectively thermodynamically destabilized, even though the bonding of AH2 is not altered. On this theme, Zuttel et al.232,233 reported hydrogen desorption properties of pure LiBH4 and 25% LiBH4 ball milled with 75% SiO2. The thermal desorptions of pure and catalyzed LiBH4 are shown in Figure 8.65. Pure LiBH4 (Figure 8.65a) exhibited a slight hydrogen desorption of 0.3 wt.% between 100 and 200 C, which is in agreement with results by Fedneva et al.229 Fusion was observed around 270 C without liberation of hydrogen. At 320 C the first

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Figure 8.65 Thermal desorption spectra of LiBH4. The gas ﬂow was measured as a function of time and the desorbed hydrogen was computed from the integrated gas ﬂow. (a) Pure LiBH4 and (b) LiBH4 mixed with SiO2 as catalyst.232

significant hydrogen desorption peak began and liberated an additional 1 wt.% of hydrogen. This first desorption transitioned to a second desorption peak starting at 400 C and reached its maximum around 500 C. The integrated amount of hydrogen desorbed up to a temperature of 600 C is 9 mass%, which corresponds exactly to half of the hydrogen in the initial compound. The end product has the nominal composition ‘LiBH’. Accordingly, the desorption equation can be written as follow: LiBH4 /LiBH2 þ H2

(R-8.14)

The thermal desorption spectra of LiBH4 mixed with SiO2 powder (25:75 mass%) shown in Figure 8.65b also exhibited three hydrogen desorption features. However, desorption started at lower temperatures, and 9 mass% of hydrogen was liberated below 400 C. The first hydrogen desorption peak started at 200 C; and the second hydrogen desorption peak started at 453 C, which includes a significant peak in desorption. The SiO2-powder catalyzed the decomposition reaction of LiBH4, and lowered the temperature for all three hydrogen desorption features. Furthermore, the pure LiBH4 sample only exhibited significant desorption above 400 C in the second hydrogen desorption peak, while the first hydrogen desorption peak starting at 200 C was the dominant peak of the catalyzed sample. The decomposition reaction of the catalyzed LiBH4 can be schematically described as follows: 3 LiBH4 /LiBH43 þ H2 T ¼ 108 C (R-8.15) 2

Hydrogen Adsorption and Storage

3 LiBH43 /}LiBH2 } þ ð1 ÞH2 T ¼ 200 C 2

(R-8.16)

1 }LiBH2 }/LiH þ B þ H2 T ¼ 453 C 2

(R-8.17)

This process can liberate up to 13.5 wt.% at the temperature around 500 C, compared with 9 wt.% at the temperature of 600 C for the pure LiBH4 sample. However, there is still total 4.5 wt.% of hydrogen remaining as LiH in the decomposition product. All attempts to synthesize LiBH4 from the elements at elevated temperature up to 650 C and 150 bar H2 pressure failed. To further destabilize the complex hydrides themselves, and thus to achieve possible candidates for hydrogen storage materials; Vajo et al.234 studied the destabilization of LiBH4 using MgH2 as a destabilizing additive. Mechanically milled mixtures of LiBH4 þ 0.5MgH2 including 2e3 mol% TiCl3 were shown to reversibly store 8e10 wt.% hydrogen. Dehydrogenation of a LiBH4 þ 0.5MgH2 mixture reacted according to: 1 1 LiBH4 þ MgH2 4LiH þ MgB2 þ 2H2 2 2

(R-8.18)

which has a maximum hydrogen capacity of 11.4 wt.%. Formation of MgB2 upon dehydrogenation stabilized the dehydrogenated state and, thereby, destabilized the LiBH4. Variation of the equilibrium pressure obtained from isotherms, which was measured at 315e400 C indicated that addition of MgH2 lowered the hydrogenation/dehydrogenation enthalpy by 25 kJ/ (mol of H2) compared to pure LiBH4. Extrapolation of the isotherm data yielded a predicted equilibrium pressure of 1 bar at approximately 225 C. However, the kinetics were too slow, reaching equilibrium took 100 hours. Aoki et al.,235 using first-principle calculations, predicted the enthalpies of dehydrogenation for LiBH4 alone and the mixture of LiBH4 þ 2LiNH2 were 75 kJ/mol H2 and 23 kJ/mol H2, respectively. Experimental examination showed that the amounts of desorbed hydrogen from LiBH4 and the mixture at 703 K and 522 K were 10.6 mass% and 7.8 mass%, respectively. Although the desorption temperature after doping decreased approximately 180 K, the dehydrogenation pressure of the mixture was much higher than that of LiBH4 alone. Au and Jurgensen236 modified LiBH4 with various additives such as TiO2, TiCl3, ZrO2, V2O3, and SnO2, and they found the hydrogen desorption initiation temperature was reduced from 673 to 473 K. The modified LiBH4 desorbed about 9 wt.% hydrogen and could be recharged to 7e9 wt.% hydrogen capacity at 873 K and 7 MPa. Although the additives reduced the dehydriding temperature and improved the reversibility, they also reduced the hydrogen storage capacity. In addition, the hydrogen storage capacity of the

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oxide-modified LiBH4 decreased gradually during hydriding/dehydriding cycling. This is possibly due to the loss of boron through the formation of BH3 released with the hydrogen stream during dehydrogenation. It has been proposed that the B loss can be prevented by selecting suitable additives, which formed intermediate boron compounds, and which changed the reaction path. In another report, Orimo et al.237 successively decomposed LiBH4 into LiH and B under 1 MPa of hydrogen at 873 K, and then rehydrided (recombined) into LiBH4 under 35MPa of hydrogen at the same temperature (873 K). The temperature of the dehydriding reaction was lowered by approximately 30 K for LiBH4 substituted (or mixed) with Mg (atomic ratio of Li:Mg ¼ 9:1) as compared to that for LiBH4 alone. This demonstrated their previously proposed method for lowering the hydrogen desorption temperature of Li-based complex hydrides. This is achieved by partial cation substitutions using smaller sized and/or higher valenced cations with larger electronegativities.238 Hydrogen storage properties of other systems such as LiBH4-Ni,239 LiBH4-CaH2,240 and LiBH4-Pt/C241 have also been investigated. The doping additive effectively decreased the dehydrogenation temperature of LiBH4; however, the irreversibility, slow kinetics, and low cycling-tolerance of the borohydrides remain to be significant challenges. From the above discussion, it is concluded that enormous success has been achieved over the last decade in hydrogen storage in complex metal hydrides. However, further efforts are still required to lower the hydrogen desorption temperature, increase the sorption kinetics, and improve cycling tolerance.

HYDROGEN STORAGE IN PURE AND REDUCED MICROPOROUS AND MESOPOROUS TI OXIDES Micro- and mesoporous Ti oxides have also been considered as hosts for hydrogen storage, due to the potential for strong interaction between hydrogen and titanium.242,243 Micro- and mesoporous Ti oxides are in many ways ideal candidates for hydrogen storage as they contain ordered channels that allow hydrogen to effectively access the interior space. The synthesis is simple, highly reproducible, and cost-effecive. Furthermore, the pore size, surface areas, and wall thickness can be systemically modified to improve hydrogen uptake.244 A unique property of this material is its ability to act as the electron acceptor, which is due to the capacity for variable oxidation states in the walls of the micro- and mesoporous structure. 245 This feature is not present in MOFs, zeolites, or porous carbon. This exceptional property, combined with the coordinative unsaturation of Ti centers, means it is likely to interact with adsorbed hydrogen more strongly than other sorbents; because binding of H2 to transition metals is strongly dependent on the electron density at the metal center and its ability to back-bond through a p-interaction into the antibonding H-H orbital, which is a Kubas-type interaction.246

Hydrogen Adsorption and Storage

Gravimetric uptake (wt.%)

(a)

5 4 3 C6-Ti C8-Ti C10-Ti C12-Ti C14-Ti

2 1 0

(b)

1.2

Gravimetric uptake (wt.%)

0

1

10

20

30 40 Pressure (atm)

50

60

70

0.8 0.6

C6-Ti C8-Ti

0.4

C10-Ti C12-Ti

0.2

C14-Ti

0 0

10

20

30

40

50

Pressure (atm)

Figure 8.66 Hydrogen storage isotherms for pristine porous titanium oxide (a) total uptake, (b) excess uptake.247

Hu et al.247 investigated the hydrogen storage properties for both the pristine and reduced microporous and mesoporous Ti oxides. Hydrogen pressureecomposition isotherms were recorded at 77 K for a series of five pristine micro- and mesoporous Ti ˚ using C6, C8, C10, C12, and C14 oxide materials with pore size ranging from 12 to 26A 244 amine templates described previously. Hydrogen total and excess uptake isotherms for this series of materials are shown in Figure 8.66.247 The microporous Ti oxide synthesized with C6 amine template (C6-Ti) showed the highest total and excess storage capacities of 5.36 and 1.08 wt.%, respectively. The corresponding total volumetric storage capacity was calculated to be 29.37 kg/m3 at 77 K and 100 atm for this sample. The general trend of storage capacity in this series of pristine titanias appeared to be one of increasing surface area, and decreasing pore size leading to more effective hydrogen storage. Both of these factors have been implicated in hydrogen

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Figure 8.67 Schematic representation of H2 binding sites in the monolayer wall of micro- and mesoporous titanium materials, where 1, 2, and 3 refer to mono, bis, and tris dihydrogen complexes, respectively, and 4 represents the compressed gas phase in the porous voids in the solid.247

physisorption to amorphous carbons and nanotubes. It is possible that s binding of H2 to Ti can occur in these microporous and mesoporous Ti oxides, where the Oh MX4L2 Ti centers245 present at the surface and in the walls are nominally electron-deficient 12 electron species possessing vacant t2g orbitals. A schematic representation of a monolayer portion of the wall is shown in Figure 8.67. For this example H2 molecules can potentially bind to the Ti centers via a side-on interaction to the d0 center through the Ti t2g set,246 or can exist as a compressed gas phase within the pore structure of the material. The pristine microporous titanium oxide (C6-Ti) was reduced by Li naphthalene (Li C6-Ti), Na naphthalene (Na C6-Ti), and bis(toluene) Ti (BTTi C6-Ti) in the liquid phase. The hydrogen sorption isotherms for these materials are shown in Figure 8.68. After reduction with bis(toluene) Ti, BTTi C6-Ti showed higher hydrogen excess gravimetric uptake of 1.14 wt.% and total volumetric uptake of 40.46 kg/m3 at 77 K and 100 atm, compared with 1.08 wt.% and 29.37 kg/m3 for C6-Ti, respectively. Considering the fact that BTTi C6-Ti had a surface area of only 208 m2/g, which is much smaller than that of C6-Ti (942 m2/g), the reduction level of Ti in the pore was a more important factor than surface area or pore volume in determining the hydrogen storage capacity in this system. The increased reduction level of the metal centers and additional Ti binding sites in the framework of the structure; which allows for easier p-back donation to the antibonding H-H orbital, and can explain the improved performance of reduced samples. Another surprising feature in these reduced materials was the unusual trend in hydrogen binding enthalpies, which showed an unprecedented increase in binding strength as the surface coverage increased (Figure 8.69).

Hydrogen Adsorption and Storage

Gravimetric uptake (wt.%)

5 C6 Ti Li C6 Ti Na C6 Ti BTTi C6 Ti

4

3

2

1

0 0

10

20

30

40

50

60

70

Pressure (atm)

Figure 8.68 High-pressure H2 isotherms for pristine and reduced C6-Ti materials at 77 K in gravimetric uptake. Solid symbols represent total uptake and open symbols denote excess uptake.247

The C6-Ti exhibited a heat of adsorption ranging from 3.0 to 4.2 kJ/mol, which decreases as surface coverage increases, this is consistent with physisorption. This usual behavior suggests the contribution of a binding mechanism, other than physisorption, possibly Kubas binding. Also, as the oxidation state of the surface titanium decreases, the heat of adsorption increases. The maximum adsorption enthalpy of 8.02 kJ/mol was obtained for BTTi C6-Ti, whose surface was most reduced as observed in the XPS studies. To further verify the existence of a Kubas-type interaction, Hu et al. investigated hydrogen storage capacities of the microporous Ti oxides reduced by the early transition metal organometallic sandwich compounds.248 Thus, when microporous titanium oxide

Heat of Adsorption (KJ/mol)

9 C6-Ti Li C6-Ti Na C6-Ti BBTi C6-Ti

8 7 6 5 4 3 2 1 0 0

0.2

0.4 0.6 0.8 H2 Sorbed (wt. %)

1

1.2

Figure 8.69 Enthalpy of H2 adsorption for pristine and Li, Na and bis(toluene) titanium reduced microporous titanium oxides.247

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Heat of Adsorption (KJ/mol)

12 Bisben Cr-Ti

10

Biscp Cr-Ti Bisben V-Ti

8

Biscp V-Ti C6-Ti

6 4 2 0

0

0.2

0.4 0.6 0.8 H2 Sorbed (wt. %)

1

1.2

Figure 8.70 Enthalpy of H2 adsorption for pristine and bis-arene and bis-cyclopentadienyl transitionmetal-reduced microporous titanium oxides.248

C6-Ti was reduced with excess bis(benzene)chromium, bis(benzene)vanadium, bis (cyclopentadienyl)chromium, or bis-(cyclopentadienyl)vanadium, the resulting materials exhibited better total volumetric storage capacities than the pristine material, ranging from 29.37 to 31.61, 33.42, 30.49, and 30.30 kg/m3 at 77 K and 100 atm, respectively. Consistent with previous reported results,247 increased enthalpy trends with increasing hydrogen surface coverage (Figure 8.70) were also observed in this series of reduced materials. These results further reflected a different mechanism of surface binding than simple physisorption, possibly involving a Kubas-type interaction. A plot of XPS Ti 3p 1/2 binding energy versus total hydrogen volumetric storage capacities for both the pristine and reduced samples is shown in Figure 8.71. Since the 45 V ol um etri c Storage (Kg/ m3)

234

40 35 30 25 20 15

BBTi C6-Ti Li C6-Ti Bisben V-Ti Biscp V-Ti

10 5

Na C6-Ti Biscp Cr-Ti Bisben Cr-Ti C6-Ti

0 34

35

36

37

38

39

Binding Energy (eV)

Figure 8.71 Plot of XPS Ti 3p 1/2 binding energy vs. the hydrogen total volumetric storage capacity for pristine and reduced samples from Hu et al.247,248

Hydrogen Adsorption and Storage

XPS binding energy can reflect the degree of reduction of the surface, an empirical linear relationship between the total volumetric storage capacity and the degree of reduction of the microframework has been established. Following the trend of Figure 8.71, it seems that even higher volumetric storage capacities are possible with more effective reducing agents. In another work, Hu and coworkers investigated the hydrogen storage capacity of Li, Na fulleride doped mesoporous titanium oxide composites.249 They reasoned that the reduced fulleride units might act as additional sites for cryogenic hydrogen binding in a mesoporous Ti oxide framework, while also creating microporous cavities within the mesoporous gallery, and enhancing absorption of hydrogen through capillary effects. The alkali fulleride doped materials showed slighter higher total volumetric storage capacities than pristine samples but a lower capacity than fulleride-free alkali metal reduced materials. This is due to the weak reducing capacity of fullerides. The impregnated fulleride units might block access to the active and coordinately unsaturated Ti metal centers, causing loss of surface area and pore volume after intercalation reaction. However, the same trend of rising enthalpies with all materials containing low valent Ti was observed. The sample reduced by Na3C60 further reduced with excess Na-naphthalene achieved the highest heat of adsorption of 6.55 kJ/ mol. These studies indicate that the surface area and pore size are not the only important parameters in determining the storage capacities. Reduction of the level of Ti in the pores is more important than surface area or pore size in this system. Another surprising feature of these reduced materials is the unusual trend in hydrogen binding enthalpies. They show an unprecedented increase in binding strength as the surface coverage increases. These features suggest a Kubas-type interaction for hydrogen binding. This was predicted by recent calculations published by researchers at National Renewable Energy Laboratory.250 However, further research is required to confirm whether this mechanism exists or not.

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9

Mercury Removal Rodolfo Monterroso, Maohong Fan, Morris Argyle Contents Introduction Mercury Species in Coal Combustion Summary of Technologies for Mercury Removal Conventional Mercury Control Technologies The Use of Air Pollution Control Devices Electrostatic Precipitation

Modeling Mercury Adsorption in An Electrostatic Precipitator Coal Cleaning and Coal Blending

Sorbent Injection in the Flue Gas Activated Carbon Injection Treated Activated Carbon Sulfur and Nitric Acid Treatment Chloride-impregnated Activated Carbon Bromine-impregnated Activated Carbon The Thief Carbon Process Mineral Oxide Sorbents Nanoscale Metal Sulﬁdes

Other Approaches for Mercury Control

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Amalgamation Mercury Oxidation Homogeneous Oxidation Heterogeneous Oxidation Catalysts for Mercury Oxidation

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Selective Catalytic Reduction Catalysts Carbon-based Catalysts Metal and Metal Oxide Catalysts SiO2eTiO2 Nanocomposite

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Induced Mercury Oxidation

Gas Phase Oxidation of Hg by Br2, Bromination Fenton Reactions used to Oxidize Mercury High-temperature Mercury Sorbents Methods for Capturing Mercury at High Temperatures from Gas Streams

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Flue Gas Components that Affect Mercury Sorption Effect of Temperature on the Mercury Sorption by ACI Sulfur-derived Components and Their Effect on Mercury Sorption Effect of H2SO4 Addition to Activated Carbon on Hg0 Removal

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Coal Gasification and Its Applications. ISBN B978-0-8155-2049-8.10009-9, doi:10.1016/B978-0-8155-2049-8.10009-9

Ó 2011 Elsevier Inc. All rights reserved.

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Effect of Oxygen in the Flue Gas Mechanism of Hg0 Removal in the Presence of SO2 by AC Summary References

286 287 288 289

INTRODUCTION Annual emissions of mercury over the world are estimated to be over 2000 tons, of which 30% are from coal-fired power plants.1,8 Although mercury is a trace element in coal with relatively low contents varying with coal rank (0.01e3.3 mg/g), its emission into the environment can be substantial due to the large energy requirements and large amounts of coal used for this purpose.1 US coal-fired power plants are estimated to currently emit approximately 48 tons of mercury per year, which accounts for about one-third of the mercury emission in the USA. On March 18, 2005, the US EPA issued the first Clean Air Mercury Rule (CAMR) for the control of mercury emissions from coal-fired power plants.2 The CAMR requires an overall average reduction in mercury emissions of about 69% by 2018.1 Mercury can go through a series of chemical transformations that convert elemental mercury to a highly toxic form.7,9 The most toxic form of mercury is methylmercury, an organic form created by a complex bacterial conversion of inorganic mercury. Methylation rates (creation of methylmercury) in ecosystems are a function of mercury availability, bacterial population, nutrient load, acidity and oxidizing conditions, sediment load, and sedimentation rates. Methylmercury enters and accumulates within the food chain, particularly in aquatic organisms such as fish and birds, causing various diseases in animals and humans.7,9 Mercury in coal is known to be associated with Fe, Cu, and S.3 In particular, it is often associated with pyritic sulfur from both pyrite and marcasite (FeS2); where it is found in the form of a solid solution and in association with the mineral sphalerite (ZnS).3,4 In addition, some Hg may also be found organically bound in coal.3 Three types of mercury to be considered are: elemental Hg(0); oxidized forms, Hg (II) or Hg(III); and particulate, Hg(p).10 In the coal combustion process, mercury is released mainly as elemental mercury, since the thermodynamic equilibrium favors this state at coal combustion temperatures. Therefore, in the combustion zone, mercury is vaporized from the coal as Hg(0). As the flue gas temperature decreases, Hg(0) is oxidized to form HgO, HgCl2, and Hg2Cl2. Concentrations of both elemental and oxidized forms range from 1 to 35 mg/m3.10 The total amount of mercury and its oxidation state depends mostly on the coal rank, since coals with a higher content in chlorine, like eastern US bituminous coal, contain high amounts of oxidized mercury, which is easier to capture (this fact

Mercury Removal

Table 9.1 Mercury distribution in coals from different regions in the United States.9 Coal-producing region Mercury content (wppm, mean) Caloriﬁc value (Btu/lb)

Appalachian, northern Appalachian, central Appalachian, southern Eastern interior Fort Union Green River Gulf Coast Pennsylvania anthracite Powder River Raton Mesa San Juan River Uinta Western interior Wind River

0.24 0.15 0.21 0.1 0.1 0.09 0.16 0.1 0.08 0.09 0.08 0.07 0.18 0.15

12,440 13,210 12,760 11,450 6360 9560 6470 12,520 8090 12,300 9610 10,810 11,420 9560

is the basis for the approach of the coal blending technique to remove mercury).9 Table 9.1 outlines the distribution of mercury in coals from different regions in the USA. Oxidized forms of mercury and particulate bound mercury are easy to capture through conventional air pollution control devices (ACPDs), such as electrostatic precipitators or fabric filters, flue gas desulfurization and scrubbers.2,6,11 Depending on the approach used and the flue gas composition, APCDs can remove up to 90% of the mercury from the combustion.13 Usually, mercury removal is much lower than 90%.13 Therefore, many studies have been conducted in order to develop sorbents and techniques that can effectively remove mercury. Mercury sorption mechanisms include: adsorption, as exhibited by activated carbon, in which van der Waals forces cause the mercury removal (although these are known as weak bonding forces, adsorbents are easy to obtain and add to most processes); chemisorption, as exhibited by fly ash when mercury oxidizers are introduced; amalgamation, as exhibited by metals like gold, silver, and copper; absorption of oxidized mercury; electrostatic precipitation of particulate mercury; and absorption via coal leaching before combustion.9e25 The current suite of mercury control technologies employs either physisorption and, or, chemisorption using sorbents like activated carbon, which is the most widely studied sorbent, Chemically treated activated carbon, unburned coal from combustion injected in the flue gas, mineral sorbents, and sulfur- and metal-containing sorbents; catalysts and chemical reactions, generally with the intent to oxidize mercury, such as plasma barrier discharges and ultraviolet radiation; scrubbing liquors; flue gas or coal additives; and combustion modifications, such as coal preheating and coal blending are also used as

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means of controling mercury dispertion.2,15 The role of fly ash and different components of the flue gas in the mercury sorption process have also been studied.12,48

MERCURY SPECIES IN COAL COMBUSTION The predominant forms of mercury in coal-fired flue gas are elemental (Hg0) and oxidized (Hg2þ, primarily as HgCl2). The thermodynamically favored species under flue gas conditions is elemental mercury.10 The scope of this section is to discuss the behavior of these species during combustion. A practical motivation for understanding the behavior of mercury during coal combustion and the mechanisms of its oxidation along the flue gas path is that mercury speciation in the flue gas determines the removal efficiency for most control technologies, since Hg2þ is more condensable and far more water-soluble than Hg0.34,36 During combustion, the mercury in the coal is transformed into three species: particle-bound mercury (Hgp), vapor-phase elemental mercury (Hg0), and vapor-phase oxidized mercury (Hg2þ), primarily in the form of HgCl2. For the best possible removal of mercury from flue gas, a high level of oxidation is helpful. Because HgCl2 is water soluble, it can be removed during the wet flue gas desalifurization process.36 Particlebound Hgp is easily removed by dust control equipment, such as baghouse filters and electrostatic precipitators. Hence, the conversion of mercury from one form to another is an important key to selecting the appropriate mercury removal technology.36 Hg0 may be oxidized to Hg2þ via homogeneous (gas phase) or heterogeneous (gasesolid) reactions.36,39 In coal, mercury starts to volatilize at temperatures below 200 C, regardless of its form. Above 600e700 C, the only stable form is elemental mercury. At temperatures below 400 C and when chlorine is present, part of the Hg0 vapor is oxidized to HgCl2(g), by direct reaction of atomic chlorine with elemental mercury. HgSO4(s) and HgO(s) become thermodynamically stable species in coal Table 9.2 Mercury species in coal.36 Mercury Temperature at maximum compound concentration ( C)

HgCl2 Hg2Cl2 HgBr2 HgS (metacinnabar, black) HgS (cinnabar, red) HgSO4 Hg2SO4 HgO

Initial temperature of appearance ( C)

Final temperature of appearance ( C)

12010 805 and 13010 1105 2055 and 2455

70 60 60 170

220 220 220 290

31010 54020 28010 5055

240 500 120 430

350 600 480 560

Mercury Removal

combustion systems at temperatures between 110 and 320 C.36 Mercury oxidation depends on the composition of the flue gas and on the amount of HCl, NOx, and SO2 present in particular. An increase in the amount of oxidized mercury has been noticed in systems that operate with selective catalytic reduction.36 Evidence suggests that fly ash has the capacity to catalyze mercury oxidation.36,39,55e58 The appearance of the different mercury species is summarized in Table 9.2. Based on thermal decomposition tests, Table 9.2 shows the temperature at which the species achieve their highest concentration, as well as the range of temperatures over which they appear.36 Based on the results of mercury speciation tests, the following facts about their equilibriums have been determined.36 • The decomposition of HgCl2 mercury phase takes place at temperatures ranging from 70 to 220 C (with maximum at 120 C).36 • The maximum decompositions of Hg2Cl2, occur at 80 C and 130 C.36 The reason for the dual maxima is a two-step reaction mechanism:36 Hg2 Cl2 4Hg0 þ HgCl2

R-9.1

HgCl2 4Hg0 þ Cl2

R-9.2

• There are two HgS crystalline structures: red HgS, or cinnabar, and black HgS, or metacinnabar.36 Metacinnabar decomposes between approximately 170 and 290 C, with two maximum values of decomposition appearing at 200 and 250 C, which suggests that its decomposition starts at 200 C, but that it decomposes completely at 250 C. Cinnabar decomposes at 310 C. The structure of cinnabar is trigonal with Hg arranged on a rhombohedral lattice, while metacinnabar’s structure is cubic. Mercury release temperatures for mercury sulfides might vary a little due to differences in crystallinity.36 • Hg2SO4 decomposes mainly at 280 C (see Figure 9.1b), which differs from the rest of the species, as experimental results show signs of a high concentration during a broad range of temperatures (120e480 C).36 • HgSO4 has a maximum decomposition value at 540 C.36 • HgO shows its maximum value of desorption at about 325 C.36 In the range 100e1600 C in chlorine-free combustion, the most abundant species is elemental mercury in vapor and gas phase with the presence of small amounts of HgO.36 The proportions of these two species vary with temperature.36 The order of the mercury appearance temperatures in matrices of fly ashes is HgCl2 < HgS < HgO.36 By inference from this data, the thermal release of mercury species is related to the vapor pressure. When the vapor pressure decreases, the mercury release temperature increases.36

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Figure 9.1 Equilibrium composition for HgSO4 and HgO, ﬁgure (a) and (b) respectively.36

According to thermodynamic equilibrium data for HgSO4 and HgO, presented in Figures 9.1a and 9.1b, respectively; when HgSO4 is heated in an inert atmosphere, it starts to decompose at 300 C and finishes decomposing at 700 C (Figure 9.1). HgSO4$HgO may also form when HgSO4 is heated, in which case it would remain stable only until 600 C (Figure 9.1a); while HgO is totally decomposed by approximately 550 C (Figure 9.1b) for both the yellow (Y) and red (R) forms.36 The difference between these two compounds is the crystallite size; although they both have the same crystallite structure, and therefore their thermal decomposition takes place at very similar temperatures (Figure 9.1b).36 In summary, thermal decomposition tests are an efficient method for identifying and quantifying mercury species from coal combustion products.36 The temperatures at the maximum decomposition rate of the mercury species can be arranged in increasing order as HgBr2 < HgCl2 < Hg2Cl2 < HgS(black) < Hg2SO4 < HgS(red) < HgO < HgSO4.36

SUMMARY OF TECHNOLOGIES FOR MERCURY REMOVAL Conventional Mercury Control Technologies The Use of Air Pollution Control Devices Existing air pollution control devices (APCDs) are capable of removing mercury to some extent, although most of them were not designed with this purpose. Existing APCDs include electrostatic precipitators (ESPs), fabric filters (FF or baghouse), flue gas desulfurization (FGD), and selective catalytic reduction (SCR).18,50 ESPs are designed to reduce fly ash emissions by creating an ionized field that removes charged particles. They have low operating costs, but a limited capacity to remove mercury because it is not generally adsorbed in the fly ash at combustion temperatures. Nevertheless, some studies suggest that when mercury is oxidized, it can be easily captured by conventional

Mercury Removal

ESPs.15,18 Baghouses are also designed to limit fly ash emissions, when flue gas is passed through a tightly woven fabric capturing particulates on the fabric by sieving. The ash collected on the filter can enhance the mercury removal. Scrubbers are installed to remove sulfur dioxide from power plant flue gas. Scrubbers use sorbents in order to remove SO2. Already installed at most coal-fired boilers to remove SO2, this device also helps to remove part of the trace metals present in the flue gas. SCR technology is used to reduce emissions of nitrogen oxides by installation of a fixed catalyst bed with reductant injection to reduce NOx to N2. In conjunction with SCR, low NOx burners create a fuel-rich primary combustion zone. This reduces the amount of thermal and fuel NOx created during combustion. SCRs also increase the amount of oxidized mercury downstream.18 Since mercury in an oxidized form is more readily captured by scrubbers, the combination of these flue gas controls may effectively capture a significant amount of mercury. One pilot study found that by installing an SCR unit, the scrubber’s mercury removal efficiency increased to about 80%.18 These efficiencies are summarized in Tables 9.3 and 9.4. Table 9.3 Efﬁciency of conventional mercury control technologies.18 Technology Hg removal (%)

Coal cleaning ESPs Baghouses FGDs SCR combined with scrubber

21 24 28 34 61

Table 9.4 Combined APCD mercury removal capabilities.13 Average mercury capture by control conﬁguration. Coal burned in pulverized-coal-ﬁred Post-combustion boiler unit emission control device Bituminous Sub-bituminous Post-combustion conﬁguration coal coal Lignite control strategy

PM control only

PM control and spray dryer adsorber PM control and wet FGD system

CS-ESP HS-ESP FF PS SDAþCS eESP SDAþFF SDAþFFþSCR PSþFGD CS-ESPþFGD HS-ESPþFGD FFþFGD

36% 9% 90% not tested not tested 98% 98% 12% 75% 49% 98%

3% 6% 72% 9% 35% 24% not tested 0% 29% 29% not tested

0% not tested not tested not tested not tested 0% not tested 33% 44% not tested not tested

CS-ESP ¼ cold-side electrostatic precipitator; HS-ESP ¼ hot-side electrostatic precipitator; FF ¼ fabric filter; PS ¼ particle scrubber; SDA ¼ spray dryer absorber system.

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Electrostatic Precipitation Electrostatic precipitators are devices that precipitate particles in a gas stream by inducing an electrostatic charge. Particles precipitate and then are filtered so that only the desired gas remains. The basic design of an electrostatic precipitator is a row of vertical wires followed by a stack of plates oriented vertically, with a usual separation of 1 to 18 cm.77 The gas stream flows perpendicular to the wires and then passes through the stack of plates, where a negative voltage of several thousand volts is applied between the wire and the plate. If this voltage is high enough, an electric discharge ionizes the gas around the electrodes. The negatively charged ions then flow to the plates and charge the particles in the gas. Ionized particles follow the electric field created by the power supply and move to the grounded plates.76,77 Particles build up on the collection plates and form a layer.77 The layer does not collapse due to electrostatic pressure, which is produced by layer resistivity, electric field, and current flowing in the collected layer. The collection efficiency of an electrostatic precipitator depends mainly on the resistivity and the particle size distribution; both properties can be determined at laboratory scale.77 A widely used model to calculate this efficiency is the Deutsch model, which assumes infinite remixing of the particles perpendicular to the stream. In this review, a modified form of the Deutsch-Anderson model is presented as the model for mercury adsorption.37

Modeling Mercury Adsorption in An Electrostatic Precipitator The Deutsch-Anderson model considers only mercury capture by suspended particles, and not capture by particulate matter already collected on ESP plate electrodes.37 It is

Figure 9.2 Electrostatic precipitator.38

Mercury Removal

assumed by the model that the sorbent is uniformly distributed all through the flue gas at the ESP inlet, that each particle gets the highest charge, and that the electric field is constant. This model uses the Deutsch-Anderson equation, which assumes initially that the suspended particle concentration varies in the flow direction, but is constant in the direction perpendicular to the electrodes. It is assumed that the particles behave as spheres of diameter dp in a log-normal distribution, therefore Equation 9.1 applies:37 ðln dp ln dpg Þ2 Eqn. 9.1 NDp ðdp Þ ¼ exp 2ln2 sg ð2pÞ1=2 dp lnsg Where NDp(dp) ¼ number density of particles of diameter dp

¼ total number of density over particles sg ¼ geometric standard deviation of the particle size distribution dpg ¼ geometric mean diameter.52 After this step, particle charge must be determined, which requires field charging and diffusion charging:37 1 Edp n ¼ 1þ2 Eqn. 9.2 þ2 4e dp kT 2p 1=2 2 n ¼ ln 1 þ dp e niN t 2e2 mi kT

Eqn. 9.3

In this case, 3 ¼ dielectric constant of the particle

n ¼ number of unit charges on a particle k ¼ Boltzman’s constant E ¼ electric field strength e ¼ the electron charge niN ¼ the concentration of ions far from the particle (where the electrostatic potential is zero) mi ¼ the particle mass t ¼ the time of exposure.52,53 After determining the charge on a particle of the specific diameter, the size-dependent terminal electrostatic drift velocity, Ues(dp), can be obtained:37 Ues ðdp Þ ¼

neECc 3pmdp

Eqn. 9.4

where C ¼ slip correction factor, which approaches 1 for micron-sized particles; m ¼ gas viscosity.

Subsequently, the Deutsch-Anderson equation determines the size-dependent, timevarying particle number density NDp(dp,t) as a result of precipitation:37

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2Ues ðdp Þ NDp ðdp ; tÞ ¼ NDp;0 ðdp Þexp t H

Eqn. 9.5

where NDp,0(dp) ¼ the number density of particles of diameter dp at the ESP inlet; H ¼ the ESP plate spacing.

Following this, the Schmidt number (Sc) and the particle Reynolds number (Re) are used to obtain the mean convective mass transfer coefficient (hm) and the Sherwood number (Sh) from the Frossling equation:37 Shd ¼

hm dp 1=2 ¼ 2 þ 0:552Red Sc 1=3 Dab

Eqn. 9.6

where Dab is the diffusivity for the Hg0eair system. The following equation gives the convective mass transfer rate of mercury to all particles of diameter dp:37 2 d _ Hg ðdp ; tÞ ¼ hm ðdp ÞNDp ðdp ÞDV •4p p rðCv ðtÞ Cs ðtÞÞ Eqn. 9.7 M 2 where Cv(t) ¼ bulk gas phase mercury concentration; Cs(t) ¼ time-dependent mercury concentration at the particle surface.

The final presentation of the same equation shows the integration of it over a time interval and over dp. After integration, Eqn. 9.7 becomes:37

Figure 9.3 Mercury removal efﬁciency by an electrostatic precipitator, model predictions compared with ﬁeld data obtained from power plants.37

Mercury Removal

MHg ðsÞ ¼

Z sZ 0

N

hm ðdp ÞNDp ðdp ÞDV •4p

0

2 dp rðCv ðtÞ Cs ðtÞÞdðdp Þdt 2

Eqn. 9.8

This last equation shows the overall particle sites, over a time interval s, convective gas-particle mass transfer within the suspension as the particle size distribution, and number density increase due to precipitation.37 This model neglects the effect of carbon fly ash in the adsorption of gas phase mercury. Figure 9.3 shows the model compared with experimental data. Lines represent the model predictions while open symbols represent the field data for full-scale sorbent injection, at the locations specified in the figures. The sorbent injection mentioned here refers to the assumption stated above for the ESP mercury capture model.37 Coal Cleaning and Coal Blending Two traditional approaches to scrub not only mercury, but also other pollutants out of coal are coal cleaning through leaching and coal blending.18 From the overall efficiencies of these two approaches, they can only reduce the relative amount of mercury emitted by a certain rank of coal compared to other ranks.18,54 That is, none of these alternatives will remove all of the mercury present in coal. Coal blending can result in a cost-effective alternative since some coal ranks have more oxidized forms of mercury.18,54 Therefore, the removal is easier when conventional air pollution control devices are available. A test was performed with a blend created by mixing a modest amount of a Western bituminous coal with a Powder River Basin (PRB) coal. Short-term operation with the blend showed that mercury emissions decreased as much as 80%.54 In addition to mercury removal, the plant benefited from the increased heat content of the bituminous coal.54 The combustors must be taken into consideration as they may be designed to burn a specific type of coal, and therefore alterations might require modifications to the whole process. Coal cleaning refers to the process of selective chemical leaching, e.g. with nitric acid.18 Many Eastern US bituminous coals are cleaned prior to use in utility power stations. This is done to reduce sulfur emissions. This coal preparation reduces sulfur content, primarily by removing pyrite from coal. In doing so, a portion of the mercury present may also be removed, as a co-benefit to sulfur reduction.18

Sorbent Injection in the Flue Gas Activated Carbon Injection Activated carbon injection (ACI) has the potential to achieve high levels of Hg removal.31,33 The performance of an activated carbon is dependent upon its physical and chemical characteristics. The physical properties of interest are surface area, pore size distribution, and particle size distribution.13 The capacity for Hg removal generally increases with increasing surface area and pore volume. Whether or not Hg and other

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Table 9.5 ACI tested in different locations.13 Test site Coal rank

ACI mercury removal (%)

PG&E NEG Brayton Point PG&E NEG Salem Harbor Wisconsin Electric Pleasant Prairie Alabama Power Gaston University of Illinois Abbot Station

94.5 94 6 25e90 73

Low-sulfur bituminous Low-Sulfur bituminous Sub-bituminous Low-sulfur bituminous High-sulfur bituminous

sorbates penetrate into the interior of a particle depends on the pore size distribution. The pores of the sorbent must be large enough to provide access to the internal surface area by Hg0 and Hg2þ, while avoiding blockage by previously adsorbed reactants. As particle sizes decrease, access to the internal surface area of the particle and adsorption rates increase. Carbon sorbent capacity is dependent on temperature, the concentration of Hg in the flue gas, the flue gas composition, and other factors.5,48 In general, the capacity for adsorbing Hg2þ will be different than that for Hg0. ACI may be used either in conjunction with existing control technologies and, or, with additional control, such as the addition of a fabric filter (FF).13 The US Department of Energy National Energy Technology Laboratory (NETL) evaluated the use of ACI in different facilities with positive results. This varied depending on the composition of the flue gas and mainly the type of coal used.13,28,29 Table 9.5 shows the values of mercury capture obtained at different locations tested. Tests were performed with a mobile sorbent injection system and a mobile test laboratory. The sorbent used is called Norit lignite-based carbon, Darco-FGD.13 This activated carbon has a bulk density of 0.51 g/cm3 and a surface area of 600 m2/g. Particle sizes range from 9 to 15 mm. The normal mercury adsorption capacity of this carbon is 200 mg Hg/g carbon. Activated carbon injection, specifically NORIT FGD activated carbon, has been studied more than any other mercury sorbent.78 Important facts to consider, while discussing its sorption mechanism, are the decrease in the equilibrium adsorption capacity when temperature increases, and when the amounts of Hg and HgCl2 increase.48,78 The effect of gas composition is also important in this case; acidic gases present in flue gas increase the sorption capacity of activated carbon, while their absence induces a low sorption capacity on activated carbon.48,59 The disadvantages of ACI include high cost and weakly interacting forces that govern the sorption mechanism. While this is not the most attractive technology, ACI is still suitable for testing at facilities that use oxygen-enriched environments for combustion or oxy-fuel combustion. Treated Activated Carbon To enhance the sorption capacity of activated carbon, some pretreatments to commercially available activated carbon have been performed to combine the advantages of activated carbon with the oxidation, amalgamation, and sorption capacities of other

Mercury Removal

compounds in order to make the material more effective while using less.19,30,32 Characteristics like BET surface area of activated carbon are generally decreased by the treatment. Nevertheless, in most cases, the sorption capacity is increased.19,20 Conventional activated carbons have proven effective for capture when applied to chlorine-containing bituminous coals.19 However, when used at plants that burn subbituminous and lignite coals, the rate of adsorption has not been as good. Prehalogenated carbon sorbents have been developed for the specific application of mercury removal at low-chlorine coal burning facilities. Tests performed by the US Department of Energy have shown more than 90% of removal with this type of sorbents, with an injection rate of 5 lb/MMacf (million actual cubic feet flue gas).19,28,30 Sulfur and Nitric Acid Treatment One of these pretreatments is obtained by mixing sulfur with activated carbon (0.3 g/g activated carbon), and then heating it to 400 C under flowing nitrogen for 6 hours.19 Another technique suggests the treatment of carbon with 6 N nitric acid for 5 hours at room temperature, followed by water wash, and 120 C dry in air. The purpose of these treatments is acidification or promotion by sulfur, which leads to oxidized forms of mercury that are more easily captured.19 Some results for these treatments are presented in Table 9.6. Chloride-impregnated Activated Carbon A study by Zeng et al. tested the use of granular activated carbon impregnated with ZnCl2.20 The activated carbon had a 0.28 mm particle size, prepared from bituminous coal by steam activation. Preparation was done by impregnation with 1 or 5% (w/v) solution for 12 hours. Impregnated carbons were then dried in an oven at 90 C. BET surface area of the 1% ZnCl2 impregnated activated carbon was 608 m2/g, while for the activated carbon impregnated with the 5% solution the surface area was 277 m2/g. The test was carried out over 8 hours to obtain a comparison of non-treated activated carbon and the two different ZnCl2 impregnated activated carbons. For the untreated activated carbon, as the temperature increased (from 50 to 200 C) the adsorption progressively decreased, as expected from typical physisorption in which the interactions between the adsorbate and the adsorbent are caused by van der Waals forces.20 The impregnation of ZnCl2 decreased both the surface area and the pore volume of the samples.20 Study of the phenomena suggests that this area was reduced due to the Table 9.6 Adsorption capacities for treated activated carbon.19 Carbon treatment Adsorption (mg Hg/g AC)

Untreated 6 N nitric acid washed Sulfur mixed and nitrogen flowed

0.078 0.141 0.096

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obstruction of internal porosity by the incorporated ZnCl2 molecules. Since the average size of pores increased with increasing zinc chloride concentration, the blocked pores were micropores, which leads to a decrease in the surface area.20 Nevertheless, the results showed a highly increased sorption capacity of these activated carbons; as shown in Figure 9.4, which shows their sorption behavior at different temperatures.20 This leads to another conclusion. Sorption capacity of activated carbon towards mercury (physisorption) increases with an increase in the surface area, but in this case the decrease led to an increase in the absorbed amount. Combining this with the evidence that micropores were occupied by ZnCl2, the sorption mechanism is deduced to be a combination of physisorption and chemisorption. The temperature dependence of the sorption mechanism confirmed this fact, since the highest value of Hg capture was achieved at 50 C, with physisorption being the dominant mechanism, because at this low temperature not many molecules are sufficiently energetic to participate in chemisorption. Therefore, the rate of chemisorption is relatively low compared to that at higher temperatures. At 100 C, physisorption decreases, since desorption takes place more readily from the activated carbon surface. At 150 C, physisorption decreases further, but there are more active adsorbate (Hg) molecules for chemisorption. Therefore, sorption is higher at 150 C compared to 100 C. At 200 C, the amount adsorbed decreases again because of the exothermic nature of the adsorption process; that is, the rate of desorption has increased to levels that negatively effect both the physisorption and chemisorption capacity of the sorbent.20,51 For this case, a mechanism for the chemisorption of elemental mercury on the Climpregnated activated carbon was proposed (Reactions 9.3 to 9.6); which states that zinc chloride was reduced by the carbon content and some Cl-containing complexes were formed.20 Then, the chloride-containing functional groups accounted for the

Figure 9.4 Activated carbon and treated activated carbon mercury sorption rates at different temperatures.20

Mercury Removal

Figure 9.5 Amount of elemental Hg adsorbed on ZnCl2 impregnated activated carbon.20

chemisorption of elemental mercury through Reactions 9.4 and 9.5. The last part of this mechanism shows that in the presence of extra Cl species, mercury tends to become four coordinated (Reaction 9.6).20 ZnCl2 þ Cn Hx Oy /Znþ ½Cl2 Cn Hx Oy

R-9.3

Hg0 þ Cl /½HgClþ þ 2e

R-9.4

Hg0 þ 2½Cl /½HgCl2 þ 2e

R-9.5

HgCl2 þ 2½Cl /½HgCl4 2

R-9.6

The results shown in Figure 9.5 were conducted at typical flue gas temperatures.20 Based on the phenomena observed, some Cl-containing complexes were inferred to have formed, which led to oxidation and sorption of the mercury.20 Bromine-impregnated Activated Carbon Hutson et al. studied mercury binding on bromine-impregnated activated carbon.23 Brominated powdered activated carbon sorbents have been shown to be quite effective for mercury capture when injected into the flue gas duct at coal-fired power plants, and are especially useful when burning Western low-chlorine sub-bituminous coals. Results indicate that the mercury, though introduced as elemental vapor, is consistently bound on the carbon in the oxidized form. The conventional and chlorinated activated carbons appeared to contain mercury bound to chlorinated sites, and possibly to sulfate species that have been incorporated onto the carbon from adsorbed SO2.23

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Figure 9.6 XPS spectra for Hg-Br reference salts and activated carbon sorbents.23

The brominated sorbents appeared to capture mercury mainly at brominated sites, confirmed by X-ray adsorption, with mainly the oxidized species of mercury bound to the sorbent.23 This indicates that chemisorption is happening due to the bromineinduced oxidation on the surface, since no elemental mercury is found on the adsorbed sites. Further study of this mechanism indicates that a subsequent binding might occur due to sulfate species. Figure 9.6 shows the X-ray photoelectron spectroscopy (XPS) data for Hg-Br reference salts and Hg-Br species on the surface of various activated carbons. The correspondence of the peaks in the spectra of both the reference and the surface species indicated that mercury was not bound in elemental form. Further reference to this sorption mechanism can be found in the study by Hutson et al.23 Another study showed that the bromination of activated carbon causes enhanced and faster Hg0 adsorption.21 The adsorption capacity increases to 0.2 mg/g C when bromination is 0.33 wt.%. Removal increases when bromine concentration increases and decreases with temperature. The sorption consumption tested at different locations showed a result of US$2,000e20,000/lb of mercury removal, while mercury removal varied from 70 to 90%.21 The Thief Carbon Process One mercury control sorbent approach that has received attention by the EPA and technology developers, is the dry sorbent injection upstream of an existing particulate device.25 In this regard, NETL and Mobotech designed a technology called the “Thief

Mercury Removal

Process”.25 This technology is very interesting because of its cost-effectiveness and efficiency. The hypothesis behind this technology is that partially combusted coal might have characteristics that are more favorable for mercury removal relative to conventional fly ash, beneficiated fly ash concentrated with unburned carbon, or other thermally/ chemically treated fly ash.25 The Thief Process forms thermally activated carbon sorbents in situ. They are obtained by inserting a lance in or near the flame, extracting a mixture of partially combusted coal and gas, and reinserting the mixture in the flue gas because its adsorptive properties are suited for mercury removal at cooler flue gas conditions. The tests conducted to date at laboratory, bench, and pilot scales demonstrate that the Thief sorbents exhibit capacities for mercury from flue gas streams that are comparable to those exhibited by commercially available activated carbons.25 For the tests, Powder River Basin (PRB) or blends of PRB/bituminous coal were burned in the pilot unit.25 The process is done with carbonaceous material continuously being withdrawn from the furnace, and then injecting it into the ductwork upstream of the existing particulate collection device. In another variation, the sorbent can be injected downstream of the plant particulate collection device, but upstream of a particulate collection device dedicated solely to the sorbent. The tests were carried out in a 500 lb/h pulverized coal-fired combustion system, nominally rated at 1.76 MW, as a reasonable simulation of the process. The setup included a combustion and environmental research facility (CERF),which is an instrumented, dry bottom, pulverized-coal combustion system that simulates the firing found in a utility power plant. The CERF is equipped with a conventional single register burner with adjustable swirl, a dual-register low-NOx burner, and options for overfire air injection and cofiring of multiple fuels with an automated process control system. There are a number of ports axially along the combustor radiant section available for the extraction of samples over a range of residence times (about 0.5 to 3 s), depending on the firing rate.24 In the CERF, unburned and burned coal is obtained from different points of the process and reinjected to the flue gas to capture mercury. Figure 9.7 shows the cross section of the CERF that was used to test this process. A similar withdrawal facility is required at a power plant to implement this technology.24,25 The surface area of the Thief sorbents can be comparable to commercially available activated carbons, ranging from 70 m2/g to 200 m2/g.24 The surface area and reactivity towards mercury are dependent upon where the solid is withdrawn from the furnace, as well as the withdrawal method employed.24 When the Thief sorbent is extracted near the burner, the content of carbon is higher, particles have a larger mean diameter and lower ash content, and, therefore, it is more favorable for this process.24 Even when activated carbon has a larger BET surface and higher sorption capacity, the cost-effectiveness of the Thief technology makes it very attractive. Activated carbon can produce a mercury removal of 94% with 1 lb/MMacf,

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Figure 9.7 Cross section of the combustion and environmental research facility (CERF).24

while the Thief sorbent is capable of obtaining 81% removal with the same sorbent/flue gas ratio.24 While initial estimated costs of installing a Thief sorbent removal system in a power plant are higher than those of activated carbon injection, ($2.89/kW for the Thief sorbent vs. $2.10 ACI calculated for a hypothetical 500 MW power plant) due to the required withdrawal facilities, the operational costs for mercury removal are considerably lower ($1,970/kg Hg removed with the Thief sorbent technology vs. $4,852/kg Hg

Figure 9.8 Hg removal by the Thief sorbent.24

Mercury Removal

removed with ACI).24 Figure 9.8 shows sorption rates for Thief sorbent using subbituminous coal compared to ACI sorption rates. The results were compared with those obtained by activated carbon and conventional fly ash.24 These results showed that the Thief sorbent had mercury sorption capacities of 0.19e1.4 mg/g at 138 C, which are similar to those of Darco FGD-activated carbon. This shows that with excellent gasesolid contact, as promoted by a packed-bed unit, unpromoted carbons exhibit a great sorption capacity. Physical adsorption appears to be involved in the sorption mechanism.24 The authors of this research state that the technology is still in development, and has not yet reached optimization and optimal costs.24 However, the technology is very interesting and feasible even for oxy-fuel combustion.24 Mineral Oxide Sorbents Research has also been performed to test noncarbonaceous materials or mineral oxides that were modified with various functional groups, to develop cost-effective sorbents for the removal of Hg(0).14 The non-carbonaceous materials were silica gel, alumina, and molecular sieves, and includes zeolites such as montmorillonite. The functional groups used to modify them were amine, amide, thiol, urea, and sodium polysulfide. Results for most of the sorbents tested indicated a negligible adsorption. Of these sorbents, the one that obtained the highest sorption value was sodium polysulfide montmorillonite, due to its high porosity and surface area, at 283 m2/g. Results indicated an adsorption of 157 mg Hg/g adsorbent at 70 C, amounting to a 93% removal efficiency, and 22 mg Hg/g adsorbent at 140 C, amounting to a 17% removal efficiency.14 Nanoscale Metal Sulﬁdes An attractive technology, due to its effectiveness, feasibility, and price; is the use of nanoscale metal sulfide sorbents.26 Although the sorption mechanism has not yet been understood, the chemical reaction mechanism includes a solid-phase oxidationreduction couple that converts mercury to mercuric ion in the lattice of cinnabar. This two-electron transfer reaction in the solid phase is a subject of current study. Of the nanoscale metal sulfides produced by this technique through the use of surfactants, only copper, gold, and silver have proven effective for mercury removal. Of these three, only copper sulfide nanoparticles are economically feasible for industrial use. Nanosized particles provide at least two times greater rates of sorption for mercury compared to microsized metal sulfides.26 Overall mercury sorption with this technology is not yet available, as quantitative results have not yet been presented, nor have prices or surface area of the sorbent been presented. Nevertheless, the surface area appears to play an important role in this mechanism, when this sorbent is compared to another consisting of the same compound but made of microparticles, the sorption rate is about two times higher.

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The concept behind the effectiveness of this compound is that copper is capable of producing an amalgam with mercury, while sulfur has oxidizing properties. Through a combination of these two capabilities, a high sorption rate can be obtained. The detailed mechanism for this sorption has not yet been studied.

Other Approaches for Mercury Control Amalgamation As introduced in the previous section, various metals can form amalgams with mercury under flue gas conditions.18 Some of them have been tested to determine their capacity for mercury removal. Gold is the most effective metal to form the necessary amalgam. In fact, the most precise mercury analyzing devices use gold to capture mercury for analysis. Some studies claim that due to the high degree of mercury capture and feasibility of separating the amalgam, gold could be reused and the mercury extracted to be sold.18 However, given that gold is an expensive precious metal, this process is not likely to be economically attractive. Silver and copper have also proven their capacity to form amalgams. These three metals can be used as sorbents in amalgamation processes. Again, for similar reasons discussed in the section on nanoscale metal sulfides, copper is the most cost-effective amalgamation candidate. Mercury Oxidation Elemental mercury can undergo homogeneous or heterogeneous oxidation.39 For temperatures below 450 C, at equilibrium, nearly all mercury should exist as Hg2þ, as discussed above. HgCl2 is generally assumed to be the dominant form of Hg2þ. Nevertheless, mercuric oxide, nitrate, sulfate, and bromide may also be formed. In flue gas, the fraction of oxidized mercury ranges from 0 to 1, which depends on the coal type and the time-temperature history of the flue gas. Based on these facts, the flue gas probably does not reach thermodynamic equilibrium. The mechanisms of mercury oxidation are not completely understood, but this section summarizes the current understanding of mercury oxidation in flue gas systems.39 Mercury oxidation in flue gas depends on, among other factors, the composition of the flue gas; while the oxidized mercury coming directly from coal depends on the factors mentioned previously. The effect of the composition of the flue gas is discussed below. Systems equipped with selective catalytic reduction increase mercury oxidation due to the elimination of NOx.39 Fly ash is believed to be able to catalyze the oxidation of elemental mercury, as will be discussed later in this section.39 Oxidation can be used as a technology for removing mercury, when it is combined with some other sorption technique. In its oxidized form, mercury is easily captured through some conventional air pollution control devices. Oxidizing mercury, as a removal technique, is often convenient in economic terms.

Mercury Removal

Homogeneous Oxidation Elemental mercury in gas phase can react with many oxidants like chlorine, HCl, chlorine radicals and ozone.39,42 The oxidation of Hg0 takes place mainly through reaction with chlorine radicals in the range 400 to 700 C, in which an abundance of chlorine radicals predominates. The Hg þ Cl reaction has low activation energy and takes place at room temperature close to the collision limit.40 The Hg þ HCl is a much slower reaction than the previous one and has high activation energy, which yields unfavorable conditions and usual operating temperatures. The subsequent reaction occurs via an intermediate product, HgCl.39 HgðgÞ þ ClðgÞ4HgClðgÞ

R-9.7

The following oxidation of HgCl is done by HCl, Cl2, or chlorine radicals.39 The amount of oxidized mercury, expressed as a fraction of Hg2þ, increases with the concentration of hydrochloric acid and the chlorine content in coal. However, the chlorine content of coal is not the decisive factor in the level of mercury oxidation, as has been seen with pilot scale coal combustion data.39 As a result, different oxidation paths must be occurring. It has been demonstrated that the homogeneous reaction between elemental mercury and chlorine is too slow to produce a considerable amount of Hg0 conversion, given that there is a big influence on the overall oxidation rate by the heterogeneous mercury oxidation on the reactor walls.39 In fact, heterogeneous oxidation is the main mechanism in the typical operation conditions of the flue gas. Besides the chlorine content in coal, the amount of oxidized mercury depends on the loss on ignition, the characteristics of the coal and any emission control devices that are located downstream, like baghouses or fabric filters.39 Heterogeneous Oxidation In the case of heterogeneous oxidation, a number of mechanisms have been proposed.39 The suggested mechanism for the heterogeneous mercury oxidation at temperatures of 300e400 C works with the following reaction and is called the Deacon process: 39 2HClðgÞ þ 1/2O2 4Cl2 þ H2 OðgÞ

R-9.8

The process could convert the large concentrations of HCl in flue gas to Cl2 in the presence of a catalyst, thereby augmenting mercury oxidation.39 Nonetheless, the equilibrium concentration of Cl2 is low and the reaction between Hg0 and Cl2 is slow. Gas-phase reactions alone are not enough to account for observed extents of mercury oxidation.43 For that reason, it is likely that a different mechanism contributes to heterogeneous mercury oxidation. The Langmuir-Hinshelwood mechanism explains the bimolecular reaction between two species adsorbed to a surface.39 AðgÞ/AðadsÞ

R-9.9

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BðgÞ4BðadsÞ

R-9.10

AðadsÞ þ BðadsÞ4ABðadsÞ

R-9.11

ABðadsÞ4ABðgÞ

R-9.12

In this case, A is Hg0 and B is possibly HCl.39 The dependence of the rate of reaction is on the concentrations of A and B, the adsorption equilibrium constant, and the rate constant for the surface reaction. Elemental mercury adsorbs to carbon in fly ash.55e58 HCl can also adsorb to carbon sites.65 Therefore, the Langmuir Hinshelwood mechanism is conceivable for catalyzing HgCl2 formation when substrates adsorb both HCl and Hg0. Many studies have mentioned the correlation between mercury oxidation and HCl concentration.39e41 Furthermore, the presence of HCl adsorbents, such as calcium oxide, have shown an inverse proportionality to the extent of mercury oxidation.39e41 SO2 competes with HCl for carbon sites on activated carbon and fly ash sorbents, and high concentrations of SO2 have been observed to inhibit mercury oxidation in this case.39, 61 Nevertheless, in different conditions and depending on the SO2 concentration, the oxidation can be enhanced or the effect can be negligible.48 More information is contained in the section dedicated to the SO2 effect on mercury sorption (see below). Schoefield suggested a mechanism for the Hg0 oxidation to HgSO4.39, 49 HgSO4 spontaneously deposits on stainless steel or platinum surfaces in flue gas containing Hg0 and SO2. The reaction is first-order with respect to mercury and zero-order to SO2. In the absence of SO2, HgO deposits. The addition of HCl to the flame after the deposit formation leads to the removal of the deposit through the reaction with HgCl2, followed by sublimation. This shows that in flue gases, HgCl2 formation follows the surface reaction to form either HgSO4 or HgO. Both species are efficiently removed from the surface through the reaction with HCl.39,49 Other explanations include the effect of NO2 and SO2 in mercury heterogeneous oxidation.39 In this model, the initiating step is the formation of Hg(HSO4)2 through reaction with SO2 on a carbon catalyst when NO2 is present, which acts as an electron sink. After which, the mercuric bisulfate reacts with NO3 to form Hg(NO3)2. Different mechanisms have been suggested to explain how this path takes place, including EleyRideal, Langmuir-Hinshelwood, and Mars-Maesen. Reactions 9.13 and 9.14 show the generic Eley-Rideal mechanism, which would represent reaction of adsorbed HCl and gas phase or weakly adsorbed Hg0 on a carbon surface. Reactions 9.15 to 9.18 show the generic Mars-Maesen mechanism.39 Eley-Rideal: AðgÞ4AðadsÞ R-9.13 AðadsÞ þ BðgÞ4A

R-9.14

Mercury Removal

Mars-Maesen: AðgÞ4AðadsÞ

R-9.15

AðadsÞ þ Mx Oy /AOðadsÞ þ Mx Oy1

R-9.16

Mx Oy1 þ 1/2O2 /Mx Oy

R-9.17

AOðadsÞ/AOðgÞ

R-9.18a

AOðadsÞ þ Mx Oy /AMx Oyþ1

R-9.18b

Nevertheless, according to current research, none of these mechanisms has been verified as the dominant mechanism for catalytic mercury oxidation.39 Some catalytic materials are believed to facilitate heterogeneous oxidation, which is typically faster than homogeneous oxidation.39 Aside from the leading mechanism for mercury oxidation, that is, heterogeneous or homogeneous, the elemental mercury oxidation in flue gas is kinetically limited.39 The following discussion lists some of the proposed catalysts for mercury oxidation. Catalysts for Mercury Oxidation Oxidation catalysts studied to date can be divided into three groups: SCR catalysts, carbon-based catalysts, and metals and metal oxides.

Selective Catalytic Reduction Catalysts Selective catalytic reduction (SCR) catalysts are used to decrease NOx concentration in flue gas.39 The catalyst is usually made of vanadium pentoxide (V2O5)/tungsten trioxide (WO3) supported on titanium dioxide (TiO2). NO is reduced by NH3, which is injected upstream of the SCR catalyst bed, at temperatures above 300 C. The V2O5 sites adsorb the NH3 strongly, while NO reacts from the gas phase or, as a weakly adsorbed species (Eley-Rideal mechanism).39 The efficiency of SCR for oxidizing mercury has also been tested for different flue gas conditions and coal types.66,67 SCR catalysts are capable of enhancing the reaction kinetics of mercury oxidation, especially when HCl is present.39 For this case, a negative correlation has been observed between the Hg0 concentration at the SCR outlet and the concentrations of chlorine (i.e. HCl) and SO2 in the flue gas at approximately 340 C.66 Increasing the NH3 concentration reduces the extent of mercury oxidation and appears to deactivate the catalyst.68 The effectiveness of several commercial SCR catalysts for mercury oxidation has been reported for power plants burning a certain mixture of coals.39,69 Under the usual

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operating conditions, Hg0 conversions are about 60e80%. The extent of mercury oxidation decreased when the NH3/NO ratio was increased, and it also decreased when the HCl concentration decreased.39,69 The mechanism for mercury oxidation on SCR catalysts is unknown, but several possibilities have been proposed. 39,43 One of them states that the oxidation of Hg0 across an SCR catalyst occurs via a similar reaction as NO reduction when HCl is present in the flue gas, with V2O5 sites adsorbing HCl and reacting with gas-phase or weakly bound Hg0.43 HCl and NH3 compete for sites on the catalyst surface, but NH3 is the dominant adsorbed species when both are present.43 The SCR catalyst can be visualized as having two different zones. In the first zone, the catalyst surface adsorbs the NH3 present and NO is reduced. The second zone occurs when there is no more NH3 present, and HCl becomes the dominant adsorbed species and mercury oxidation takes place. Some other mechanisms have been proposed by Niksa and Fujiwara, Eswaran and Strenger, and Senior.43,44,57 Senior’s model predicts accurately the results at laboratory and plant scale, taking into account the most important parameters for such oxidation.57 Nevertheless, at this point, it is impossible to determine the reaction mechanism from the available data.39 Further research is required to determine the nature of this reaction.

Carbon-based Catalysts Activated carbons are general sorbents that can remove a number of different species from flue gas. Furthermore, activated carbon injection (ACI) is an established method for adsorbing elemental and oxidized mercury from combustion flue gas, as discussed in prior sections. In this section, its function as a catalyst for mercury oxidation will be discussed. The presence of adsorbed Hg0 and HCl opens the possibility for heterogeneous oxidation of mercury.39 Therefore, carbon-containing materials can be used as mercury oxidation catalysts. Studies have shown that Hg0 will also adsorb to the carbon content of fly ash, and that the extent of mercury oxidation in flue gas depends upon the amount of unburned carbon present in fly ash and the coal type.55e58 Most of the research in carbon-based catalysts has focused on fly ash.39 When mercury oxidation occurs on fly ash particles, it is assumed that it takes place in carbon sites of the fly ash.59 Thus, there is a relationship between unburned carbon and the extent of mercury oxidation, since the unburned carbon can be estimated from the amount of fly ash present in the flue gas and the amount of carbon present in the fly ash, although it usually turns out to be an overestimation.39,58 There are many factors in play when trying to understand the catalytic effect of carbon in fly ash on mercury oxidation. In the presence of fly ash, when the concentration of HCl in flue gas is higher, there is an increase in Hg2þ.39,64 Also, if the coal type used contains high amounts of chlorine, greater formation of Hg2þ is favored.39,60

Mercury Removal

There is a significant interaction between fly ash and NOx, perhaps controlled by the NO/NO2 ratio.39,61 NO2 can heterogeneously oxidize Hg0, though this reaction is considered of less importance compared to chlorination. Moreover, NO2 enhances the extent of mercury oxidation, while NO inhibits oxidation. Mechanisms describing this effect are uncertain.39,61 The role of SO2 is unclear with respect to fly ash catalyzed oxidation. The effect described in section 4 for SO2 may be used to understand its role in this case. SO2 can inhibit Hg0 oxidation.39,56 Several bituminous and sub-bituminous fly ashes can convert up to 20e50% of the Hg0 in a simulated flue gas to Hg2þ.39,62 As some tests have shown, lignite-derived fly ashes exhibit less catalytic ability, and that most of them will oxidize less than 10% of Hg0. Results have shown that the effectiveness of the fly ash as a catalyst decreases with time, which is dependent upon on the coal type from which the fly ash originated.39,62 For example, for lignite-derived flue gas, fly ash and carbon catalyst mercury conversion fell from 100% to 0% after 18 weeks of exposure. In the presence of sub-bituminous derived flue gas, one carbon catalyst deactivated completely within 2000 hours. Another carbon catalyst fell from converting 100% of the incident Hg0 to 80%.39,63 Thief carbon is an alternative to fly ash as a catalyst for mercury oxidation.39 Thief carbon is produced by taking partially combusted coal out of the furnace. Thief carbon has a high weight percentage of carbon (30e50%) and works as an effective sorbent for Hg0. 24,25,39 Due to this high content of unburned carbon, Thief carbon has catalytic properties.39 Thief carbon can oxidize up to 75% of the Hg0 in a bench-scale slipstream of flue gas at 140 C.39 The reaction mechanism for mercury oxidation on fly ash is unknown.39 Among the possible mechanisms, the following are listed: a LangmuirHinshelwood reaction between HCl and adsorbed Hg0, or an Eley-Rideal reaction between gas-phase HCl and adsorbed Hg0. Since fly ash contains metal oxides and chlorides, a Mars-Maessen reaction is also possible. The fact that the Hg0 adsorption onto the fly ash surface is temperature dependent, suggests that the process is a physical adsorption.39 Nevertheless it has also been proposed that a chemical sorption takes place in this process, and that the temperature dependence is part of a pre-equilibrium step of the mercury sorption.43 Apart from of the sorption mechanism, it has been noticed that adsorbed Hg0 is included in the oxidation reaction.39 Further research must be performed in order to fully understand this mechanism.

Metal and Metal Oxide Catalysts Besides V2O5 present in SCR catalysts, a number of other metals and metal oxides have been studied as potential mercury oxidation catalysts.39 As with the SCR and carbonbased catalysts presented before, the reaction mechanisms for the metal and metal oxide catalysts are unknown. A logical assumption is that the possible reaction mechanisms noted above, Langmuir-Hinshelwood, Eley-Rideal, and Mars-Maessen; and the

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Deacon process, which is used to produce Cl2 exist as possibilities for these catalysts as well.39 Iron and its oxides may catalyze Hg0 oxidation.70 The extent of mercury oxidation in the presence of fly ash increased with the Fe3O4 content of the ash. Also, flue gas containing HCl and fly ash containing Fe2O3 achieved 90% oxidation of the incoming Hg0.71 Iron catalysts have been tested at both laboratory and pilot scale.40 Results have shown less than 50% Hg0 oxidation in lab tests of an iron catalyst held in a fixed bed at 150 C.40 Nevertheless, a different study achieved 30% mercury oxidation with flue gas derived from sub-bituminous coal and 90% oxidation with flue gas derived from bituminous coal; the catalyst was composed of coldside iron.63,72 Other studies suggest that iron and iron compounds present in stainless steel might make it an effective catalyst.73,74 Noble metals, including copper, gold, silver, and palladium, have been tested as potential mercury oxidation catalysts.39 Ghorishi et al. tested a model fly ash containing CuO that oxidized >90% of the Hg0 present in simulated flue gas.71 The same study found that CuCl present in a model fly ash is reactive enough to oxidize Hg0 even in the absence of gas-phase HCl.71 Gold, silver, platinum, copper, and mixtures of these metals have also been proposed as potential catalysts.39 A test of a gold catalyst achieved >95% oxidation of the Hg0 present in a simulated flue gas at 70 C.74 Hg0 oxidation across a gold catalyst in the presence of Cl2 has been reported to produce values of up to 40e60% conversion.75 Studies of palladium catalysts showed less than 30% mercury oxidation.62 Tests of cold-side palladium catalysts burning lignite, sub-bituminous, and bituminous coals achieved >90% oxidation.63 Palladium maintains its catalytic capacity for periods of up to 10 months.39 The palladium catalyst can also be regenerated to nearly new performance by purging with either N2 or CO2.72 Due to its high activity over a long period of time and its ability to be regenerated, palladium is a very good option as a catalyst. Furthermore, its effect can be considered when designing a sorbent. Other suggested catalyst materials include iridium, manganese, manganese oxide, iridium, and iridium/platinum.39

SiO2eTiO2 Nanocomposite A novel catalytic method involves the use of TiO2 and ultraviolet (UV) radiation.45 In the presence of UV light, Hg0, and water react on the surface to form a TiO2$HgO complex. This reaction is first-order with respect to Hg0 concentration and is capable of capturing 99% of Hg0 at low temperatures. Nevertheless, at temperatures above 110 C, the oxidation decreases because of mercury desorption from the catalyst surface due to the higher temperature.45 Under ultraviolet (UV) irradiation, hydroxyl (OH) radicals can be generated on the surface of TiO2 and then oxidize Hg0 into mercury oxide (HgO), which is retained on

Mercury Removal

the particle surface due to its low vapor pressure45. The reaction mechanism is described by Reactions 9.19 to 9.23 45: TiO2 þ hv/e þ hþ

R-9.19

H2 O þ Hþ /OH

R-9.20

hþ þ OH /$OH

R-9.21

hþ þ H2 O/$OH þ Hþ

R-9.22

$OH þ Hg0 /HgO

R-9.23

where e and hþ represent an electron-hole pair on the TiO2 surface. Removal of Hg0 has been achieved using photocatalysts in the form of either in-situ generated TiO2 particles, or a high surface area silica gel doped with TiO2 nanoparticles (SiO2eTiO2 nanocomposite). The nanocomposite is beneficial due to its high surface area and open structure, this allows irradiation by UV light and minimizes mass-transfer resistance for Hg0. The efficiency of Hg0 removal using a SiO2eTiO2 nanocomposite reached 99% at low relative humidity and room temperature.45 However, studies have shown that at high water vapor concentrations, mercury sorption is decreased by the adsorption of water vapor.45 The decrease is proportional to the water vapor fraction present in the flue gas. Typically, the water vapor fraction in the flue gas accounts for 6 to 12% by volume. For that reason, water vapor may have a great inhibitory effect on Hg0 removal in flue gas over these catalysts. Nevertheless, this catalyst has been designed for application at the cold end of the boiler convective pass, where the usual temperature of the flue gas is in the range 120 to 150 C.45 As noted before, coal-derived flue gas contains various trace gas components such as HCl, SO2, and NOx. Their concentrations are dependent on the type of coal burned, coal firing systems, boiler operating conditions, and control processes. These trace gases are relevant to heterogeneous adsorption and mercury oxidation on catalysts, activated carbons, or fly ash under flue gas conditions. For this reason, the performance of the SiO2eTiO2 nanocomposite has been tested in two simulated flue gases.45 The first flue gas, containing 8% water, 30 ppm HCl, 1200 ppm SO2, 300 ppm NO, and 10 ppm NO2, represents flue gas from high rank (bituminous) coals that contain higher chlorine and sulfur content. The second flue gas, which contains 12% water, 10 ppm HCl, 400 ppm SO2, 300 ppm NO, and 10 ppm NO2, represents flue gas from low rank, sub-bituminous and lignite, coal. This type of coal contains less chlorine and sulfur, but more moisture.45 The effect of NO in the first flue gas inhibited the promotional effects of HCl and SO2. Hg removal in the second flue gas was less than in the first, due to the higher concentration of H2O, and lower concentrations of HCl and SO2.

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Figure 9.9 Oxidation efﬁciency of the SiO2eTiO2 nanocomposite and effects of the ﬂue gas.45

Therefore, high rank coals are preferable to low rank coals for the application of the nanocomposite. Minimizing the effect of NO could substantially help to improve the catalytic effect and removal efficiency. Figure 9.9 summarizes the efficiency of the nanocomposite for mercury oxidation and the effects of the flue gas composition. The flue gas composition affects significantly the extent of the mercury conversion. For idealized conditions in this study, the SiO2eTiO2 nanocomposite can capture elevated amounts of mercury.45

Mercury Removal

In summary, the flue gas component effects were that HCl and SO2 promoted Hg oxidation and capture, while H2O and NO inhibited Hg removal, and the effect of NO2 did not prove significant.45 Induced Mercury Oxidation Two notable methods studied for induced mercury oxidation are bromination and Fenton reactions. Therefore, they are proposed as possible technologies for the mercury removal in this study.

Gas Phase Oxidation of Hg by Br2, Bromination Bromination is the process in which elemental mercury is oxidized by bromine gas. The resulting product is mercury bromide or HgBr2.12 This process is based on the principle that mercury is a large atom, with 80 electrons, and therefore is very polarizable or “soft”. Since bromine is a large nonpolar molecule, the London interaction forces of HgeBr2 should be greater than those exhibited by chlorine. Therefore, a higher capacity of oxidation among halogens is expected from bromine gas.12 In order to determine the effect of the oxidation of mercury by the bromine gas in a flue gas environment, all of the components present must be taken into account to determine the real effect of the reaction of bromine with mercury.12 The effect of different gases present in the flue gas was studied. This study showed that most gases caused no significant effect on the flue gas; the exception was the gas NO.12 The effect of carbon fly ash was also studied, which proved that it plays an important role. The reaction mechanism indicates that the determining step of this process is the chemisorption of the elemental mercury on the fly ash adsorbed bromine, because the oxidation rate constant increased with the use of carbon fly ash. The reaction rate was second order, 6.00.51017 cm3 molecules1 s1 at 298 K, which is 2 orders of magnitude larger than that exhibited by Cl2. This supports the idea of a stronger van der Waals interaction for this case.12 The kinetics for this study were performed using a large excess of oxidants.12 In such conditions, the half life of mercury is independent of the initial Hg0 concentration used. UV light exposure had negligible effect on the mercury oxidation rate. The Gibbs free energy of the reaction was calculated to be 35.6 kcal/mol at standard conditions and 30.8 kcal/mol at 500 K, indicating a thermodynamically favorable condition for the reaction to occur. The disappearance of Hg0 showed an exponential decay, which indicated a pseudo-first-order reaction with respect to Hg0.12 The total oxidized mercury without carbon fly ash was 50% with 52 ppm Br2.12 When the effect of fly ash was studied, the total amount of oxidized mercury was 60% with an injection of 0.4 ppm of gaseous bromine in the flue gas. The positive effect of fly ash in the removal decreased with increasing temperatures; this could lead to possible lower results on a plant scale.12 To determine the effect of the fly ash, the experiments

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Figure 9.10 Mercury oxidized through bromination, expressed in terms of elemental mercury left after a certain exposure time.12

were carried out in flasks coated with halocarbon wax to minimize the effects of light and to impregnate the wax with fly ash. Figure 9.10 shows mercury oxidation in different conditions of bromine, exposure to UV light (in order to understand its effect which resulted practically negligible), and fly ash.12 Since this experiment was done in batch conditions to establish reaction order, it is important to consider differences in a continuous process at a power plant, particularly the residence time of the flue gas exposed to bromine. The presence of NO in the flue gas had an effect in the gas phase oxidation of Hg0 in the presence of bromine.12 NO promotes the oxidation rate of Hg0 at NO concentrations lower than 8 ppm, but inhibits the rate at concentrations greater than 10 ppm. The proposed mechanism to explain this phenomenon includes the following reactions:12 Hg0 þ Br2 4HgBr2

R-9.24

HgBr2 þ M4HgBr2 þ M

R-9.25

HgBr2 þ NO4Hg0 þ Br2 þ NO

R-9.26

NO þ Br2 4NOBr2

R-9.27

Hg0 þ NOBr2 4HgBr2 þ NO

R-9.28

Hg0 and Br2 have large electron clouds and, therefore, are highly polarizable. They attract to each other by London forces, which results in the formation HgBr2).12 This excited species then relaxes to form the stable HgBr2 molecule after colliding with a third molecule M, such as N2. However, the excited molecule of mercury bromide could separate back into Hg0 and Br2 upon electrostatic interaction with NO, which has an unpaired electron. In the same way, the NOBr2) van der Waals cluster can form as

Mercury Removal

a result of the interaction of the polarizable Br2 with NO. The collision of Hg0 with NOBr2) gives a different path for the production of mercury bromide. Reactions 9.25 and 9.28, which form excited states of molecules, are assumed to be in fast equilibrium. Using the steady-state hypothesis, the rate of formation of the oxidized species of bromine (HgBr2) is the following:12 d½HgBr2 k24 k25 ½MHg0 ½Br2 k27 k28 ½Hg0 ½Br2 ½NO ¼ þ dt k24 þ k25 ½M þ k26 ½NO k27 þ k28 ½Hg0

Eqn. 9.9

When the concentration of NO is small enough, to the point where k-24þk25[M] [ k26[NO], the influence of the concentration of NO on the first term is negligible. While the second term shows an increase in the formation rate of HgBr2 when the concentration of NO increases.12 As [NO] increases, then a reduction in the formation rate of HgBr2 can be expected in the first term of Equation 9.9. When the magnitude of that reduction overcomes the increase of the second term, the overall formation rate of HgBr2 slowly decreases. The most important conclusion from this mechanism is that under most flue gas conditions, NO is expected to have concentrations large enough to inhibit the oxidation of Hg0. This effect has also been observed with NO2, assuming that the odd number of electrons is the main reason why the effect is shown; as other molecules with an even number of electrons did not show the same effect, i.e. SO2, CO, and O2.12 There are several possible disadvantages of bromination; all of which require additional research. They include further oxidation of the bromine (although this is unlikely), the conditions required for actual coal combustion or gasification facilities, such as the residence time of the flue gas necessary to achieve a desired overall mercury oxidation, and bromine residues in the process which can emit another toxic gas.

Fenton Reactions used to Oxidize Mercury This procedure is based on the conversion of Hg(0) to Hg(II) in coal-derived flue gas based on the well-known Fenton reactions.10 Once oxidation is complete, a mercury control strategy can be implemented in a wet scrubber. Tests for this process were done at laboratory and bench scale.10,27 Fenton reactions refer to the promoted oxidation reaction of hydrogen peroxide (H2O2) by the presence of the ferrous ion. This principle is used in this case to oxidize the elemental mercury. The reaction mechanism has not been proven, although some studies propose suitable models.10 According to models published for the oxidation of elemental mercury,35,46 the reaction sequence is the following: Fe3þ þ H2 O2 4Fe2þ þ $OOH þ Hþ

R-9.30

Fe2þ þ H2 O2 4Fe3þ þ $OH þ OH

R-9.31

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2Hg0 þ 2$OOH þ 2Hþ 4Hg2 2þ þ 2H2 O2

R-9.32

Hg2 2þ þ 2$OOH þ 2Hþ 4Hg2þ þ 2H2 O2

R-9.33

The key step of this reaction chain in the oxidation of elemental mercury is the OOH radical, which is generated from the reduction of the ferric ion by hydrogen peroxide.10 The bench scale results showed that Hg(0) oxidation can be achieved by the Fenton reactions.10 The oxidation rate is quantitatively dependent on the residence time of the Hg stream in the solution. An average of 75% oxidation of Hg(0) was achieved. This type of reaction can be accomplished by the use of different salts and an oxidizing molecule like hydrogen peroxide. For this test, iron-based components obtained the highest efficiency, specifically Fe2(SO4)3 and FeCl3. The pH of the solution (where the sorbent is located) had a significant effect on the oxidation of mercury; values in the range of 1.0 and 3.0 had the best effect on the solution. The results obtained show effectiveness, reasonable cost, and the possibility of obtaining harmless by-products. The tests were also conducted in oxygen-enriched combustion, which is of special interest in this review. The results were practically the same as those obtained with an air-blown combustion; which indicates that this technology can be used with oxy-fuel combustion, although it does not present any specific advantage for it.10 A pilot scale test was also conducted using this principle.27 The oxidation results were not as high as those obtained in the bench scale tests, but the procedure is feasible and easily adjustable to a power plant. The setup used was a vertical combustor research facility that included an ESP, a baghouse, and a condensing heat exchanger (CHX). The primary mercury scrubber was the CHX. The procedure was to oxidize the mercury and then capture it in the CHX. Only the ferric solutions from the bench scale test were used, which were ferric sulfate and ferric chloride, and three different types of coal: bituminous, sub-bituminous, and lignite. During the process, H2O2 was reinjected, while there was no need to increase the initial amount of the ferrous compounds since they act only as catalysts. The results varied significantly with the coal types due to of the mercury species produced by each of them; which is determined by the amount of chlorine present. The coals with a higher concentration of chlorine, like US eastern bituminous, present more oxidized mercury in the flue gas, and, therefore, the sorption was higher. For the higher coal rank, total removal was 81%. For a Canadian western sub-bituminous coal, the overall removal was 45%. For Saskatchewan lignite coal, the average removal was only 40%.27

High-temperature Mercury Sorbents In a coal gasification process, where syngas temperatures are considerably higher than those of the flue gas in coal combustion, the use of physical adsorption might not be

Mercury Removal

capable of adsorbing high amounts of mercury. For physisorption, when temperature rises, desorption rates increase, and the sorption capacity decreases. Carbon-based sorbents are unsuited for mercury capture at gas temperatures higher than 205 C.47 Some sorbents have been developed to capture mercury at higher temperatures. In order to accomplish this goal, it is necessary to change the mechanism to chemical sorption and amalgamation, which can also be performed at a high temperature. The main disadvantage for this type of sorbents is high cost. Research is currently being carried on to obtain cost-effective sorbents.47 For example, Granite and Pennline tested a series of sorbents that can be used at typical syngas temperatures.47 Methods for Capturing Mercury at High Temperatures from Gas Streams Recent research has been performed to obtain sorbents that are capable of operating over a wide range of temperatures. A patent was filed by Granite and Pennline 47 for hightemperature mercury sorbents. This patent mentions that a feature of the tested sorbents was the use of the active metal sorbents’ exterior surface, first at a lower temperature and then allowing the metal to adsorb mercury at a second higher temperature in the interior of the sorbent.47 The sorption temperature ranged between 175 C and 370 C.47 The sorption mechanism claimed is amalgamation or a chemical reaction. An advantage of this temperature sorption range is that the gasifying process maintains its thermal efficiency. The surface areas of the tested sorbents are between approximately 1 m2/g and 1000 m2/g, which increases as temperatures increase. Such sorbents increase the capacity with increasing temperature in the range 150e290 C. The sorbents used are alloyed to enhance their resistance to corrosion.47 These sorbents can be regenerated when heated above 925 C, or by contact with an acid.47 Both regeneration processes are one-step. Therefore, the process is reversible and allows for proper disposal or use of the captured mercury. Different metals, including iridium, palladium, platinum, and ruthenium, as well as their alloys, were considered for the creation of such sorbents. The iridium-platinum alloy appears to be effective for both flue and syngases.47 Another advantage of using high temperatures for mercury sorption, in addition to the possibility of using this process in gasification, is the ability to use warm gas cleanup methods. These methods already exist for scrubbing hydrogen sulfide from the effluent gases.47 All of these characteristics make such a process viable for IGCC systems, municipal waste incinerators, hazardous waste incinerators, and any type of coal-fired power plant. However, the cost of using this sorption system can be very high due to the used of noble metals. Therefore, for further validation of such processes at an actual coalburning, or similar process locations where mercury is emitted, current prices must be evaluated. This should occur even when regeneration is possible, due to the high cost of metals. This study cites breakthrough characteristics of the sorbent as the resistance to

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corrosion of many gases such as molecular chlorine, hydrogen chloride, sulfur dioxide, and sulfur trioxide. But it is important to consider the economic viability before any technology is implemented. Another notable characteristic of these sorbents is their capacity to remove other species beside mercury, like cadmium, arsenic, and selenium.47 The mechanism of mercury capture is a two-step process, which includes surface adsorption and the formation of a solid solution, which is an amalgam formation by mercury and the metal present in the sorbent.47 A metal film is created in the substrate surface, which helps both the amalgamation and the adsorption. At temperatures of 120e150 C, mercury is trapped in the region of the active metal; therefore, the main mechanism of the capture is surface adsorption. A subsequent diffusion of mercury into the inner regions of the metal increases with temperatures above 150 C. Such temperatures enable the noble metal to initiate the amalgamation process. This process is commenced by the expansion and opening of a metal substrate’s surface due to the heating. Mercury adsorbed first at a lower temperature can penetrate into the metals’ inner regions at a second higher temperature.47 Figures 9.11 and 9.12 depict the process and mechanism of adsorption. Figure 9.11 depicts the adsorption-desorption apparatus for metals onto the tested metal sorbents.47 Figure 9.12 depicts the metal sorbent on a substrate’s surface. In Figure 9.11, a gas stream (number 2), which contains mercury, is fed into the process. Subsequently, it interacts with the metal sorbent (5) while residing in a controlled atmosphere. This atmosphere is defined by housing (4), so temperature and pressure control are possible. In this way, the interior of the housing (4) serves as the reaction zone for the process. In the optimal case, the housing should include heating for the reaction zone (6). Following contact with the metal sorbent, treated gas (7) can be treated for remaining pollutants in (8). Finally, the treated gas stream (9) is released through any further processing steps. The process provides for regeneration of the sorbent by heating

Figure 9.11 Process ﬂow diagram for high-temperature mercury sorbents.47

Mercury Removal

Figure 9.12 Sketch of a high-temperature mercury sorbent particle.47

at temperatures above the adsorbing-reaction temperatures. The desorbed mercury is afterwards processed for use or proper disposal.47 As mentioned before, the diffusion of mercury into the metal increases with temperatures above those used for the initial adsorption step, this is due to the amalgamation mechanism.47 These temperatures open the lattices of the active, hard, highmelting metals tested, which contain the metal film to make their interiors accessible for increased mercury capture. The expansion and opening of a metal substrate’s surface upon heating is shown as numeral 12 in Figure 9.12. This diffusivity of mercury into the sorbent lattice structure provokes the formation of the solid solution mentioned above. Therefore, Hg is adsorbed onto the surface at a given temperature and penetrates it at a second higher temperature.47 The metals tested (iridium, palladium, platinum, and ruthenium) were chosen due to their mechanical properties that are more suitable for the conditions used in this process, since their melting points are 2447 C, 1554 C, 1772 C, and 2310 C, respectively.47 Ir is corrosion-resistant, hard, elastic metal that has an elastic modulus of 570 GPa.47 The rest of the metals are corrosion-resistant as well, and their respective elastic moduli are 121 GPa, 168 GPa, and 220 GPa, respectively. All of these metals have a hexagonal or cubic close-packed lattice structure. Mercury only adsorbs on the surface of these metals at ambient temperatures, most likely due to their hardness.47

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Table 9.7 Mercury capture from simulated ﬂue gas by Ir, Pa, Pt and Ru.47 Metal Mercury capture at 290 C (%)

Pt Pd Ir Ru

100 76 3 1

Table 9.8 Mercury sorption capacity for the Iridium sorbent at different temperatures.47 Sorbent Capacity (mg/g) Temperature ( C)

Ir Ir Ir

2.11 2.69 2.9

204 232 260

The experiments were carried out at NETL. For the tests, a gas concentration of 270 ppb Hg was used.47 For each trial, the sorbent was 10 mg of metal sorbent in a vertical packed-bed reactor; the gas was 36% mol CO, 27% mol H2, and the balance of nitrogen. The flow rate was 60 ml/min. Final results are shown in Tables 9.7 and 9.8, cold vapor atomic absorption spectroscopy was used to measure the amount of mercury.47 The iridium sorbent was measured alone to determine its overall sorption capacity.47 The results presented in Table 9.8, as well as other experimental data, show that in the range 260 C to 290 C the highest values of mercury removal for these metal sorbents are obtained.47

FLUE GAS COMPONENTS THAT AFFECT MERCURY SORPTION The composition and conditions of the flue gas can significantly affect the adsorption of elemental mercury on carbon-based sorbents as well as on any other sorbents. Nevertheless, the effect has mainly been studied on the former sorbents; a discussion is presented here on how this composition can potentially affect carbon-based sorbents.48 Sulfur oxides, HCl, nitrogen oxides, water, and oxygen have been reported to affect this sorption.48 The data in Figure 9.13 show that the presence of some gases actually affects the sorption behavior of mercury with activated carbon, specifically O2, H2O, and SO2.48 A more detailed explanation of the effect of each gas will be presented in this section.

Effect of Temperature on the Mercury Sorption by ACI Physical adsorption decreases with increasing temperature.48 This can also be observed in the statistical mechanical derivations of physical adsorption models like the Langmuirian

Mercury Removal

Figure 9.13 The combined presence of O2, SO2, and H2O in ﬂue gas increases mercury removal by activated carbon.48

model.79 No detailed derivation of this model is presented here as it is not within the scope of this review, but the following two equations are included to highlight the inverse proportionality between temperature and physical adsorption:79 q ¼

N x P ¼ ¼ Ns 1þx P þ P0

kB T 2pkB T 3=2 P0 ¼ zL h2

Eqn. 9.10

Eqn. 9.11

P0 stands for the partial pressure of the adsorbates and q stands for the surface sites covered by the adsorbate.79 The rest of the parameters are defined elsewhere, but the relevance of these two equations is that the effect of increasing system temperature increases the partial pressure of mercury, which results in a reduction of the total adsorption. This decrease will stop once both pressures equalize.79 The effect of temperature has been noted in almost every study done on mercury removal through physical adsorption. Figure 9.14 shows the sorption capacity of activated carbon at three different temperatures, 60 C, 80 C, and 100 C.48 The same relationship between temperature and chemical sorption cannot be made. In many cases the chemisorption rate increases with temperature due to the temperature dependence of the reaction processes. In this case, the temperature dependence is more complex, since it varies from case to case. For a detailed understanding of a particular chemisorption and its temperature dependence, it is recommended to review the behavior of the sorbent of interest in each case; as no general correlation is established due to the complexity of the systems, especially when both physical and chemical sorption mechanisms are involved. See Zeng et al.20 as an example.

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Figure 9.14 Effect of temperature on Hg0 removal by activated carbon.48

Sulfur-derived Components and Their Effect on Mercury Sorption Hg0 can be removed in the presence of SO2 and absence of H2S by ACI.48 However, the mechanism and role of these sulfur components are not well understood.48 SO2 and O2 are necessary for mercury removal.48 Moreover, the presence of water accelerates the sorption.48 Figure 9.15 shows Hg0 removal by fly ash activated carbon in the presence and absence of SO2. This demonstrates that the sorption capacity of a given activated carbon changes when the flue gas composition is varied. In this particular case, SO2 increases this adsorption by orders of magnitude.48 From the observed phenomena, the following reaction mechanism is proposed to contribute to elemental mercury removal, acting as intermediate step: SO2 þ 1/2O2 4SO3

R-9.34

SO3 þ H2 O/H2 SO4

R-9.35

Figure 9.15 Effect of the presence of SO2, on the Hg0 removal by ﬂy ash activated carbon.48

Mercury Removal

The relationship between these components and the mercury removal will be more clearly demonstrated in the next section.

Effect of H2SO4 Addition to Activated Carbon on Hg0 Removal Following the assumption that H2SO4 formed from O2, H2O, and SO2 contributes to Hg0 removal; sulfuric acid addition to activated carbon has been shown to increase mercury sorption capacity.48 From this result, if Reactions 9.34 and 9.35 occur over the activated carbon, the deposited H2SO4 is assumed to contribute to mercury removal. For this reason, the effect of treatment of the AC with SO2, O2, and H2O has also been studied. The contribution of Reactions 9.34 and 9.35 was confirmed. However, the Hg0 removal rate is considerably higher when activated carbon is treated as described above and SO2 is not present in the flue gas. The explanation for this is that H2SO4 helps the adsorption by being adsorbed on the activated carbon and not necessarily just by being present in the flue gas; it also causes an enhanced adsorption potential and enthalpy of adsorption. Also, a new narrowly distributed microporosity is formed, along with an increased surface polarity due to the impregnation.48 Besides the SO2 reduction, the presence of other sulfur oxides (SOx) has been reported to inhibit the mercury removal.48 In summary, SO2 must be present in the flue gas for activated carbon to capture elemental mercury; but when the activated carbon has already been treated with sulfuric acid, the absence of SO2 enhances the adsorption rate. Figure 9.16 shows the effect of the different gases present in the flue gas and the sorption behavior of sulfuric acid treated activated carbon when conditions are changed.48 Figure 9.16a shows the mercury removal by activated carbon treated with sulfuric acid in the presence and absence of SO2, while Figure 9.16b shows the removal of mercury without SO2 in the simulated flue gas.48 These figures demonstrate that sulfuric

Figure 9.16 Mercury removal in the presence or absence of SO2 using activated carbon treated with sulfuric acid (a) and activated carbon without sulfuric acid (b).48

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acid can improve the mercury removal for carbon-based sorbents when it is impregnated on the sorbent’s surface and that SO2 must be present to enable the mercury sorption, except if the activated carbon sorbent is already treated with H2SO4. Nonetheless, mercury adsorption by activated carbon in presence of SO2 is not directly correlated to the rate of formation of H2SO4 (rather the effect of sulfuric acid on Hg sorption is more related to the aforementioned reasons). This was determined by conducting sorption experiments with injection and suppression of SO2. From the observations, a reaction mechanism was proposed to explain the phenomenon that occurs after the first two reactions (9.34 and 9.35):48 Hg þ 1/2O2 4HgO

R-9.36

HgO3 þ H2 SO4 4HgSO4 þ H2 O

R-9.37

The gas phase oxidation of mercury by oxygen happens mostly at 300 C, but it also occurs below this temperature at a slower but appreciable rate when activated carbon is present since it works as a catalyst.39,48

Effect of Oxygen in the Flue Gas Based on the previously discussed results, the effect of oxygen in the flue gas is also important for the removal of mercury. Therefore, the effect of O2 has been tested when SO2 is not present.48 The results confirmed that oxygen accelerated the adsorption.48 However, the Hg0 removal efficiency remained high even in the absence of O2 in the feed gas. Figure 9.17 shows the results of these experiments involving oxygen content of the flue gas. The oxygen concentrations tested were 0 and 5% v/v.48 It has been assumed from the results that some small amount of oxygen remained in the system after displacement with oxygen free gases, so the remaining oxygen could contribute to the reaction.48

Figure 9.17 Effect of O2 on the elemental mercury adsorption in the absence of SO2.48

Mercury Removal

Mechanism of Hg0 Removal in the Presence of SO2 by AC Based on the experimental results of the study of the effects of different components in the flue gas,48 the mechanism shown in Figure 9.18 for mercury removal when SO2 is present was proposed. Figure 9.18 illustrates that SO2 reacts with O2 (route 1) and then forms H2SO4 (route 2). Hg reacts with oxygen and produces HgO (route 3).48 Then, HgO reacts with H2SO4 and produces a sulfate-containing complex. However, some of these compounds, like the HgSO4-containing complex, and the formation of another complex from the reaction of mercury and sulfuric acid, have not yet been confirmed. The suppression of mercury removal by the presence of SO2 can be explained if routes 5 and 6 occur. Following the assumptions and the experimental results, the mercury compounds on the surface include HgO and the Hg2þ SO2 4 complex, and possibly contain HgSO4. Consequently, the reduced Hg compound comes from the reduction of HgO by SO2 and the remainder is Hg2þ/SO4 complex. This reducibility was also studied. From the results, the following reaction was assumed to take place:48 HgO þ SO2 4Hg0 þ SO3

R-9.38

Reaction 9.39 is affected by the presence of O2, due to the occurence of Reaction 9.3449. It was observed that Reaction 9.39 occurs predominantly over Reaction 9.34, which was determined in the absence of O2 when the Hg0 evolution from HgO was examined. It was also observed in the case of the evolution of Hg by SO2 treatment. The occurrence of Reaction 9.34 in the SO2-N2 system has also been directly confirmed.49 However, the evolution of Hg was not observed. The necessity of the presence of H2O for the evolution of Hg from HgO reduction with SO2 is still unknown. The following reaction possibly occurs during Hg evolution:49

Figure 9.18 Schematic diagram of the adsorption of elemental mercury in the presence of SO2.48

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HgO þ H2 O þ SO2 4Hg0 þ H2 SO4

R-9.39

The amount of HgO compared to the amount of elemental mercury is controlled by the oxidation of Hg0 by O2 to HgO, as shown by Reaction 9.36; and by the reduction of HgO to Hg0 by SO2 (Reaction 9.38).48 The removal rate of Hg0 by oxidation decreases with increasing temperature, as mentioned above and shown in Figure 9.14, due to the reduction of HgO with SO2 occurs preferably compared to Hg0 oxidation.48

SUMMARY To date, many technologies have been proposed to enhance mercury removal with costeffective results. Existing controls already possess the capability for capturing mercury, Table 9.9 Removal capacities of different mercury capture techniques.

Process

Coal cleaning ESPs Baghouses FGDs SCR combined with scrubber Coal blending Regular activated carbon injection Norit Darco FGD activated carbon

Mercury removal (%): maximum achieved value at the conditions presented in each study.

21 24 28 34 61 80 65 to 94.5 >90

Amount of Hg adsorbed (mg Hg/g sorbent): maximum achieved value at the conditions presented in each study.

78e200 mg/g (in absence of HCl) 200 mg/g (in absence of HCl) to 1610 mg/g (with HCl present) 141 mg/g 96 mg/g

6 N nitric acid washed Sulfur mixed and nitrogen flowed 5% ZnCl2 activated carbon

N/A N/A

Bromine impregnated activated carbon Thief sorbents Bromination Fenton reactions Sodium polysulfide montmorillonite Nanoscale metal sulfides Iridium

70e90

800e900 mg/g (after 8 h of exposure) N/A

81 N/A 40e81 17 (93 achieved at 70 C)

190e1380 mg/g N/A N/A 22 mg/g

N/A 96

N/A 2,900 mg/g

N/A

Mercury Removal

although in most of these cases, these controls might not meet the regulations to be implemented in the near future. Therefore, not only are combinations of available processes needed to meet requirements, but it is also important to implement a technology capable of withstanding the different variables exhibited by different coal combustion processes, including different types of coal, air- or oxygen-enriched combustion, and different control devices available at the facility. Understanding mercury sorption mechanisms is an important current study. Physical adsorption as well as chemical adsorption can effectively capture mercury; other chemical reactions, like the oxidation of mercury, can also assist this process. Halogens and other oxidizers, due to their oxidizing properties towards mercury, have been the subject of study in this regard. Metals and metal compounds, due to their amalgamating properties, have also shown effectiveness in mercury capture. Table 9.9 summarizes the different approaches and their mercury removal capability. Further research is still necessary to obtain sorbents that are more cost-effective. Some combinations of sorbents can also be used both for research and for direct application, e.g. oxidizers in combination with APCDs like ESPs or scrubbers, or oxidizers combined with adsorbents like activated carbon or Thief sorbent; which represents a low-cost solution that could potentially enhance its overall mercury removal capacity when combined with another technology. Therefore, tests of different combinations of the technologies described in this review are yet to be studied and carried out at actual coal-burning facilities. This will one day determine which removal option yields the best solution.

REFERENCES 1. US EPA. Clean Air Mercury Rule, <www.epa.gov/mercury>. March 15, 2005. 2. Feeley TJ, Murphy JT, Hoffmann JW, et al. DOE/NETL’s Mercury Control Technology Research Program for Coal-Fired Power Plants. EM, 2003:16-23. 3. Davidson M, Clarke LB. Trace elements in coal. IEA Perspective Report IEAPR/21; 1996. 4. Sloss LL. Mercury emissions and effects: the role of coal. London: IEAPER/ 19, IEA Coal Research; 1995. p. 39. 5. Granite EJ, Presto AA. Comment on the Role of SO2 for Elemental Mercury Removal from Coal Combustion Flue Gas by Activated Carbon. Energy & Fuels. 2008;22:3557-3558. 6. Feeley TJ, Brickett LA, O’Palko BA, et al. Field testing of mercury control technologies for coal-fired power plants. DOE/NETL Mercury R&D Program Review, May; 2005. 7. Office of Air Quality Planning and Standards and Office of Research and Development. Mercury Study Report to Congress. Washington DC: US Environmental Protection Agency, ; 1997. 8. Global Mercury Assessment. United Nations Environment Programme (UNEP) Chemicals. Geneva, Switzerland: UNEP, ; 2002. 9. Tewalt SJ, Bragg LJ, Finkelman R. Mercury in U.S. coaldabundance, distribution, and modes of occurrence, . USGS Fact Sheet FS-095-01 September (2001). 10. Lu D, Anthony EJ, Tan Y, et al. Mercury removal from coal combustion by Fenton Reactions Science Direct. Fuel 2007;86:2789-2797.

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11. Granite EJ, Pennline HW, Hargis RA. Novel sorbents for mercury removal from flue gas. Ind Eng Chem Res. 2000;39:1020-1029. 12. Shou-Heng L, Qiangyan N, Rong Liu Z, et al. Using bromine gas to enhance mercury removal from flue gas of coal-fired power plants. Environ Sci Technol. 2007;41:1405-1412. 13. Control of Mercury Emissions from Coal-fired Electric Utility Boilers. Air Pollution Prevention and Control Division. U.S. Environmental Protection Agency. 14. Joo YL, Yuhong J, Timc K, et al. Development of cost-effective noncarbon sorbents fo Hg0 removal from coal fired power plants. Environ Sci Technol. 2006;40:2714-2720. 15. Brown TD, Smith DN, Hargis RA, et al. Mercury measurement and its control: what we know, have learned, and need to further investigate. JAWMA, 1999;49:628-640. 16. Ochs T, Patrick B,Oryschchyn D, et al. Proof of Concept for Integrating Oxy-fuel Combustion and the Removal of all Pollutants from a Coal Fired Flame, Jupiter oxygen Corporation. USDoE/Albany Research Center. 17. Gerdemann S, Summers C, Oryschchyn D, et al. Developments in integrated pollutant removal for low-emission oxy fuel combustion. U.S. Department of Energy/ARC-2005-026 18. Mercury Control Options for Coal-fired Power Plants, Clean Air Network Fact Sheet. www.cleanair.net. 1999 August. 19. Removal of mercury from coal derived synthesis gas, Western Research Institute, 2006 U.S.D.O.E. 20. Zeng H, Feng J, Guo J. Removal of elemental mercury from coal combustion flue gas by chlorideimpregnated activated carbon, Science Direct. Fuel 2004;83:143-146. 21. Sun W, Yan N, Jia J. Removal of elemental mercury in flue gas by brominated activated carbon. Zhongguo Huanjing Kexue (China Environ Sci.) 2006;26:257-261. 22. Apogee Scientific Inc. Assessment of low cost novel sorbents for coal-fired power plant mercury control, final report, March 2004, 23. Hutson N, Atwood B, Scheckel K. XAS and XPS Characterization of Mercury Binding on Brominated Activated Carbon. Environ Sci Technol. 2007;41:1747-1752. 24. O’Dowd WJ, Pennline HW, Freeman MC, et al. A technique to control mercury from flue gas: The Thief Process, Fuel Process. Technol 2006;87:1071-1084. 25. Granite EJ, O’Dowd WJ, Pennline HW, et al. The Thief process for mercury removal from flue gas, U.S.D.o.E. National Energy Technology Laboratory. 26. Peterson E, Martellaro P, Moore G. Nanoscale metal sulfides for mercury sorption. Am Chem Soc Div Fuel Chem. 2008;53(1):85-86. 27. Tan Y, Lu. D, Anthony E, et al. Mercury removal from coal combustion by Fenton reactions: Pilot scale tests, Science Direct. Fuel 2007;86:2798-2805. 28. Amrhein G, Sjostrom S. Long term carbon injection field test for 90% mercury removal in a PRB unit with an SCR spray dryer and fabric filter, ADA Environmental Solutions. 29. Machalek T., Full-scale activated carbon injection for mercury control in flue gas derived from North Dakota lignite, Presented at the A&WMA/EPA/DoE/EPRI Combined Power Plant Air Pollutant Control Mega Symposium, Washington DC, August 30eSeptember 2, 2004. 30. Thompson J. Enhancing carbon reactivity for mercury control: field test results from Leland olds, Presented at the A&WMA/EPA/DOE/EPRI Combined Power Plant Air Pollutant Control Mega Symposium, Washington DC, August 30eSeptember 2, 2004. 31. Hargis RA, O’Dowd WJ, and Pennline HW. Pilot-scale research at NETL on sorbent injection for mercury control, Proceedings of the 26th International Technical Conference on Coal Utilization and Fuel Systems, March 2001. 32. Enhancing carbon reactivity in mercury control in lignite-fired systems. Quarterly Technical Report to DOE under Cooperative Agreement No. DE-FC26-03NT41989. Energy & Environmental Research Center, University of North Dakota, ; February 2005.

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33. Machalek T, Richardson C, Dombrowski K. Full-scale activated carbon injection for mercury control in flue gas derived from North Dakota lignite, pubs/MEGA 41989 stanton.pdf, accessed October 2006. 34. Schroeder K, Cardon C, Hesbach P, et al. NETL in-house characterization of mercury in coal combustion by-products, in Proceedings of Mercury Control Technology R&D Program Review Meeting, Pittsburgh, PA, 2005 July 12e14. 35. Kobayashi T. Oxidation of metallic mercury in aqueous solution by hydrogen peroxide and chlorine. J Jpn Soc Air Pollut 2008;22(3):230-236. 36. Lopez-Anton MA, Yuan Y, Perry R, et al. Analysis of mercury species present during coal combustion by thermal desorption. Fuel, 2010;89:629-634. 37. Clack H. Mercury capture within coal-fired power plant electrostatic precipitators: Model evaluation. Environ Sci Technol. 2009;43:1460-1466. 38. 39. Presto A AA, Granite EJ. Survey of catalysts for oxidation of mercury in flue gas. Environ Sci Technol 2006;40(18):5601-5609. 40. Hargrove O, Carey T, Richardson TC, et al. Enhanced control of mercury and other HAPs by innovative modifications to wet FGD processes, Report to DoE/FETC. U.S. Department of Energy Agreement No. DEAC22e95PC95260; Radian International; 1997. 41. Laudal D, Brown T, Nott B. Effects of flue gas constituents on mercury speciation. Fuel Process Technol 2000;65e66:157. 42. Sliger R, Kramlich J, Marinov N. Towards the development of a chemical kinetic model for the homogeneous oxidation of mercury by chlorine species. Fuel Process Technol. 2000;65e66:423. 43. Niksa S, Fujiwara N. A predictive mechanism for mercury oxidation on selective catalytic reduction catalysts under coalderived flue gas. J Air Waste Manage Assoc 2005;55:1866. 44. Eswaran S, Stenger H. Understanding mercury conversion in selective catalytic reduction (SCR) catalysts. Energy Fuels 2005;19:2328. 45. Li Y, Murphy P, Chang-Yu. Removal of elemental mercury from simulated coal-combustion flue gas using a SiO2eTiO2 nanocomposite, Fuel Process. Technol. 2008;89:567-573. 46. Walling C, Goosen AJ. Mechanisms of the ferric ion catalyzed decomposition of hydrogen peroxide. Effect of organic substrates. Am Chem Soc. 1973;95:2987-2991. 47. Granite EJ, Pennline HW. Method for high temperature mercury capture from gas streams, United States Patent No. 7,033,419 B1. 48. Azhar Uddin M, Yamada T, Ryota O, et al. Role of SO2 for Elemental mercury removal from coal combustion. flue gas by activated carbon,. Energy & Fuels 2008;22:2284-2289. 49. Schofield K. Let them eat fish: Hold the mercury. Chem Phys Lett. 2004;386:65-69. 50. Benson S, Laumb J, Crocker C, et al. SCR catalyst performance in flue gases derived from subbituminous and lignite coals. Fuel Process Technol. 2005;86:577. 51. Bansal RC, Donnet JB, Stoeckli F. Active Carbon. Marcel Dekker; 1988:504. 52. Friedlander SK. Smoke, Dust, and Haze: Fundamentals of Aerosol Dynamics. Oxford University Press; 2000. 53. White HJ. Industrial Electrostatic Precipitation. Addison-Wesley; 1963. 54. Sjostrom S. Evaluation of Sorbent Injection for Mercury Control. ADA Environmental Solutions for U.S. DOE Cooperative Agreement No. DE-FC26e03NT41986 Topical Report No. 41986R11, ; 2006. 55. Hassett D, Eylands K. Mercury capture on coal combustion fly ash. Fuel 1999;78:243. 56. Serre S, Silcox G. Adsorption of elemental mercury on the residual carbon in coal fly ash. Ind Eng Chem Res. 2000;39:1723. 57. Senior C, Johnson S. Impact of carbon-in-ash on mercury removal across particulate control devices in coal-fired power plants. Energy Fuels 2005;19:859. 58. Gibb W, Clarke F, Mehta A. The fate of coal mercury during combustion, Fuel Process. Technol. 2000;65e66:365.

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59. Olson E, Miller S, Sharma R, et al. Catalytic effects of carbon sorbents for mercury capture. J Hazard Mater 2000;74:61. 60. Kellie S, Cao Y, Duan Y, et al. Factors affecting mercury speciation in a 100-MWcoal-fired boiler with low-NOx burners. Energy Fuels 2005;19:800. 61. Laudal D, Brown T, Nott B. Effects of flue gas constituents on mercury speciation, Fuel Process. Technol. 2000;65e66:157. 62. Hargrove O, Carey T, Richardson C, et al. Enhanced control of mercury and other HAPs by innovative modifications to wet FGD processes, report to DoE/FETC. US Department of Energy Agreement No. DEAC22e95PC95260; Radian International; 1997. 63. Blythe G, Miller S, Richardson C, Searcy K. Enhanced control of mercury by wet flue gas desulfurization systems site 2 results, Report to US DoE/NETL. US Department of Energy Agreement No. DEAC22e95PC95260; Radian International; 2000. 64. Galbreath K, Zygarlicke C. Mercury transformations in coal combustion flue gas, Fuel Process. Technol. 2000;vol. 65e66:289. 65. Mangun C, Benak K, Economy J, et al. Surface chemistry, pore sizes, and adsorption properties of activated carbon fibers and precursors treated with ammonia. Carbon, 2001;39:809. 66. Laudal D, Pavlish J, Galbreath K, et al. Pilot-scale evaluation of the impact of selective catalytic reduction for NOx on merury speciation, Report to U.S. DoE/NETL. Grand Forks: US Department of Energy Agreement No. DEFC26-98FT40321; Energy and Environmental Research Center, University of North Dakota; 2000. 67. Blythe G. Pilot Testing of Mercury Oxidation Catalysts for Upstream of Wet FGD Systems, Quarterly Technical Progress, Report to U.S. DoE/NETL. U.S. Department of EnergyAgreement No DE-FC26e01NT41185; URS Corporation; 2003. 68. Machalek T, Ramavajjala M, Richardson M, et al. Pilot evaluation of flue gas mercury reactions across an SCR unit, In Proceedings of the DOE-EPRI USEPA-AWMA Combined Power Plant Air Pollutant Control Symposiums. The MEGA Symposium; 2003. 69. Senior C. Oxidation of mercury across SCR catalysts in coal-fired power plants burning low rank fuels, Final Report to DOE/NETL. US Department of Energy Agreement No. DEFC26e03NT41728; Reaction Engineering International; 2004. 70. Dunham G, DeWall R, Senior C. Fixed-bed studies of the interactions between mercury and coal combustion fly ash, Fuel Process. Technol. 2003;82:197. 71. Ghorishi S, Lee C, Jozewicz W, et al. Effects of fly ash transition metal content and flue gas HCl/SO2 ratio on mercury speciation in waste combustion. Environ Eng Sci 2005;22:221. 72. Blythe G, Marsh B, Miller S, et al. Enhanced control of mercury by wet flue gas desulfurizations site 3 topical report, Report to US DOE/NETL. US Department of Energy Agreement No. DE-AC22e95PC95260; URS Corporation; 2001. 73. Granite E, Pennline H. Catalysts for oxidation of mercury in flue gas. U.S. Patent Application; 2005. 74. Meischen S, Van Pelt V. Method to control mercury emissions from exhaust gases. U.S. Patent 2000;6:136-281. 75. Zhao Y, Mann M, Pavlish J, et al. Application of gold catalyst for mercury oxidation by chlorine. Environ Sci Technol. 2006;40:1603. 76. IUPAC. Compendium of Chemical Terminology, 2nd edn. (the “Gold Book”). Compiled by A.D. McNaught and A. Wilkinson. Blackwell Scientific Publications (1997). XML on-line corrected version: 2006- created by M. Nic, J. Jirat, and B. Kosata; updates compiled by A. Jenkins. 77. 78. Yang H, Xua Z, Fan M, et al. Adsorbents for capturing mercury in coal-fired boiler flue gas. J Hazard Mater 2007;146:1-11. 79. Langmuir I. The constitution and fundamental properties of solids and liquids. J Am Chem Soc. 1916;38:2221-2295.

CHAPTER

10

CO2 Sorption Soonchul Kwon, Maohong Fan, Herbert F.M. DaCosta, Armistead G. Russell, Kathryn A. Berchtold, Manvendra K. Dubey Contents Introduction Sorption-based CO2 Separation Processes Physical Adsorption Pressure Swing Adsorption (PSA) Vacuum Swing Adsorption (VSA)

Chemical Adsorption Adsorbent Performance for CO2 Capture Physical Properties and Capture Capacities of Different Sorbents Zeolites Zeolite Selectivity Purity and Recovery of Zeolite The Cost of CO2 Separation Using PSA and VSA with Zeolite Equilibrium and Kinetic Adsorption

293 296 297 297 299

300 301 302 303 305 306 307 309

Activated Carbon Amine-modiﬁed Adsorbents Organic/Inorganic Hybrid Materials Lithium Zirconate Mineral-based Adsorbents

311 314 318 322 325

Thermodynamics and Kinetics Kinetic Improvements

326 327

Sodium-based Sorbents Hydrotalcite-like Compounds (HTlcs) CO2 Capture Development Strategies References

328 331 333 334

INTRODUCTION Carbon dioxide is naturally released from the respiration of animals, plants, and microorganisms and is consumed by photosynthesis and oceanic disolution.1 These large natural sources and sinks are in quasi balance. However, anthropogenic emissions from fossil fuel combustion are cause of concern because they have perturbed this natural carbon balance, increasing levels of atmospheric CO2 by 40% over preindustrial levels to about 390 ppm today (IPCC). This rise has been linked to observed warming and ocean acidification, and will grow rapidly to meet the world’s growing energy demand, Coal Gasification and Its Applications. ISBN B978-0-8155-2049-8.10010-5, doi:10.1016/B978-0-8155-2049-8.10010-5

Ó 2011 Elsevier Inc. All rights reserved.

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Figure 10.1 Anthropogenic emission sources of carbon dioxide worldwide in 2006.2

particularly in the developing world. Among green house gases, carbon dioxide accounts for 94% of the total greenhouse gas emissions, and over 80% of the anthropogenic CO2 emissions in the world is generated through fossil energy production (Figure 10.1).1 Since carbon dioxide contributes the largest portion of the overall greenhouse gas emissions, its emission reduction has been the focus of efforts to mitigate climate change.3 From a breakdown of the sources of worldwide carbon dioxide emissions in 2007 by various sectors (Figure 10.2), power generation is the largest source of CO2 emissions, accounting for about 41% of the total. Consequently, CO2 separation and sequestration

Figure 10.2 Carbon dioxide emissions from fossil fuel combustion as a function of usage sectors in 2007. )Other contains commercial/public services and agriculture/forestry.2

CO2 Sorption

from the flue gas of power plants can play a very important role in reducing global CO2 emissions. Typically, three technical options are considered for the reduction of CO2 emissions: (1) decreased energy consumption and more efficient energy usage; (2) development of renewable energy sources and non-fossil fuels such as hydrogen; and (3) development of capture and sequestration technologies to trap more CO2 underground or in the ocean. Among these options, capture and sequestration (CCS) has the potential to contribute to global emission decreases sooner and faster than the development of renewable energy sources.4 Briefly, for CCS technologies, carbon dioxide is captured and separated from the flue gas of power plants, and then piped for storage in geological reservoirs. The separation of CO2 from the captured flue gas is essential to avoid storing a large amount of N2. Typical flue gases of power plants contain about 72% N2, 8e12% CO2, 8e10% water vapor, and smaller concentrations of other pollutant species such as SOx and NOx, these flue gasses are pre-treated by scrubbers to improve the purity of inlet CO2 gas.5,6 The main purpose of carbon sequestration is to build a CCS system in a fossil-fuel-fired power plant, which is considered as a point emission source. This will improve CO2 reduction efficiency and to store the pure stream CO2 generated. This will prevent CO2 emission into the atmosphere. This is critical because fossil fuels are still the principal energy source that is affordable, and available amounts of fossil fuels are large enough to meet the world’s demand. CCS has been already used for the oil and gas industry. However, the major concern with separation methods for CCS is the increased energy consumption and cost resulting from the process of temperature swing, or pressure swing to initiate CO2 desorption. Furthermore, the low separation efficiency necessitates alternative or improved CO2 separation methods that have low energy demands, high CO2 sorption capacity, and good regeneration capacity.7 CO2 capture capacity is defined as the amount of CO2 captured per unit mass of adsorbent used (e.g. g CO2/g sorbent). Regeneration capacity denotes how much capture capacity is recovered after multi-cycle sorption. Among the various processes, high selectivity adsorption has gained attention because of its potential for simple operation, low corrosiveness,4,8 and overall low cost.9,10 In particular, in terms of environmental issues regarding the use of chemicals, molecular sieves and molecular baskets of solid adsorbents are environmentally friendly. The high-potential adsorbents should have (1) high selectivity and adsorption capacity for CO2, (2) fast adsorption and desorption rates, (3) high stability to be cycled multiple times, (4) a low attrition rate, (5) low energy requirements for the regeneration of CO2 after adsorption, and (6) should be scalable to megaton/year CO2 capture rates. In particular, for practical power plant applications, rapid adsorption of CO2 from the flue gas is highly desirable for high efficiencies of the multi-cycle adsorption/desorption process, leading to lower quantities of adsorbent materials being used, and smaller control device volumes being required. Since the adsorption capabilities of sorbents depend on

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the CO2 concentration in the power plant stack, temperature, pressure, affinity of sorbent to CO2, surface area, pore size and structure of sorbents, the improvement of systems and adsorbents requires a process level understanding of the dependence on these variables.11e13 This chapter provides an overview of the currently studied CO2 adsorption technologies with various solid adsorbents. We evaluate their performance as well as the process equilibrium and kinetics when they are available in the literature. Future trends and strategies for CO2 separation are also explored.

SORPTION-BASED CO2 SEPARATION PROCESSES Aqueous and dry processes have different CO2 sorption characteristics. In aqueous sorption, CO2 is dissolved into the solvent. In solid sorption, CO2 molecules are chemisorbed to the sorbents by strong chemical bonds, or physisorbed by weaker intermolecular bonds resulting from the interactions between sorbent and CO2 molecules. Since many solvents are degraded by reactive compounds such as fly ash, SOx, and NOx, aqueous sorption should proceed after additional pretreatments are done to remove the toxins. These processes should include electrostatic precipitation and flue gas desulfurization. Initially, aqueous processes dissolve CO2, and the concentrated CO2 solution is transported to a regeneration column where the CO2 is separated from the solution. The solvent is then regenerated and recycled to a new inlet of flue gas. This solution process leads to high capital and operating costs due to its high energy consumption for the solvent regeneration; its high rate of corrosion of the process equipment; its fast rate of evaporation causing solvent losses; and its high rate of degradation in the presence of oxygen.14e16 On the other hand, dry reactions without solvent have simpler design and operation due to the absence of pre-treatment process; the requirement for small pressure changes only for physical adsorption; and the requirement for the increased temperature only in the regeneration column for chemical desorption. In some cases, the reaction rate in the solidegas phase can be quite fast even at typical flue gas temperatures and pressures while the reaction kinetics in the aqueous phase tends to be slower. In addition, due to the temperature limitations and low CO2 solubility in water, the aqueous systems tend to require high pressures. Furthermore, in order to improve the aqueous processes, additional chemicals (such as hydrochloric acid and acetic acid) are required in the dewatering stage, adding to process cost and impairing durability (due to corrosion). Thus, for a highly efficient CO2 capture system with cost-effectiveness, adsorption using solid adsorbents is mainly discussed. For the evaluation of the CO2 sorption activity with respect to the separation process, two principal modes of adsorption, physical and chemical, are discussed. Later, currently researched adsorbents will be introduced.

CO2 Sorption

Physical Adsorption Physical adsorption, physisorption, of molecules involves relatively weak intermolecular forces. These forces include dispersion, dipolar or Van der Waal interactions between the adsorbent surface and the adsorbate CO2 molecules. Such interactions operate at long range (i.e. large intermolecular distances) but have weak strength, with the energy release being of the order of magnitude of the enthalpy of condensation. Physisorption does not provide a large redistribution of electron densities at either the adsorbent surface or the adsorbate. Its energy can be adsorbed as lattice vibrations and dissipated as thermal motion. The physisorption enthalpy of adsorption is typically of the order of 20 kJ/ mol.17 In terms of kinetics, adsorption rates are relatively fast because the process is not associated with large activation energy barriers. Furthermore, the energy needed to regenerate CO2 from the physisorbed state is typically low and hence attractive. However, in some cases, mass-transfer controlled processes could inhibit van der Waals interactions. For physisorption, the two applicable options, pressure and vacuum swing adsorption, are discussed below. Pressure Swing Adsorption (PSA) Pressure swing adsorption (PSA) has been applied to separate gas mixtures, such as carbon dioxide capture in ammonia production and in hydrogen purification.6,13,18e32 In terms of the high cost-effectiveness, PSA is viewed as an attractive approach due to its simple operation, high performance at ambient temperatures, high regeneration rate, and low energy intensity.33 In the PSA process, gas species can be separated from a gas mixture at high pressure and low temperature, by being passed through a reactor containing the sorbent. The pressure is then reduced, releasing CO2 from the adsorbent surface relatively easily since electrons between adsorbent and adsorbate are not shared. Figures 10.3 and 10.4a illustrate examples of pressure swing adsorption systems. PSA operation includes:13,29 (1) pressurization of the inlet gas (over 3 atm), (2) adsorption with inlet gas at high pressure, (3) depressurization to atmospheric pressure, releasing CO2 at the bottom of desorption column, and (4) desorption CO2 gas from the adsorbent with purging gas at ~1 atm. After flue gases are introduced into one chamber column 1 is pressurized, resulting in the CO2 adsorption. The applied pressure is transferred to column 2. When column 2 is pressurized, column 1 is depressurized to ~1 atm and carbon dioxide is separated from the flue gas. For the desorption steps, the inlet CO2 gas stream is stopped and N2 is only introduced to desorb CO2 after depressurization. The cycle continues as a switching mode from adsorption column to desorption column. The mechanism of the PSA system is based on either equilibrium thermodynamics selectivity or kinetics selectivity. Thermodynamic equilibrium selectivity depends on the different gas concentrations at equilibrium state of the gas molecular mixture for the separation processes. For PSA using equilibrium selectivity, strongly adsorptive gas

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Figure 10.3 Schematic diagram of pressure swing adsorption system.34

components remain in the adsorbent column, whereas weakly adsorbed species are discharged in the high-pressure gas streams. Kinetic selectivity, on the other hand, is based upon having different diffusion rates of gas molecules in the non-equilibrium system. Faster diffusing gas species go to the sorbent column while the slower diffusing gases flow out. Thus, the diffusion rate determines the selectivity mechanism. For the optimization of the PSA process, operating conditions and system factors are determined by the rate-determining mechanism. Operating processes have evolved to improve the efficiency of PSA separations. To optimize the bed size and sorbent usage, different adsorptive sorbents such as activated carbon and zeolites, which have strong CO2 affinity, are utilized within the layered sorbent bed where activated carbon sorbents are placed ahead of a zeolite sorbents.20 In the first layered column, activated carbon mainly adsorbs CO2 species and zeolites in the second layered column capture the residual CO2 components. This approach was introduced in the 1970s; but it was not well suited for industrial facilities due to the many limiting factors, including the low sorption capacity of adsorbents and unstable operating controls. PSA has three significant drawbacks. First of all, the short cycle times bring high “switch” losses, which are the losses of feed gas in the bed after venting out in the depressurization step.35 In addition, the short cycle times can cause the rate change of inlet flow, leading to the unstable pressure in the column during the plant operation. Secondly, PSA processes typically add impurities to the gas components at low pressures and these components will adsorb more strongly than carbon dioxide. Thirdly, PSA has a low cost-effectiveness for CO2 recovery from flue gas steams. CO2 recovery costs can lower when the CO2 concentration in a flue gas stream is high. However, separation processes operate at thousands of kPa and cannot separate CO2 to less than about 10 kPa

CO2 Sorption

Figure 10.4 Operational steps of pressure and vacuum swing adsorption systems.5,13 (a) Pressure swing adsorption system. (b) Vacuum swing adsorption system.

partial pressure due to the low CO2 partial pressures. Thus, compression systems are used to reduce high pressure to the normal operating pressure, but this additional energy demand causes extra costs. In spite of these weaknesses, PSA still plays a significant role in the hybrid system due to low cost and simple operation. Vacuum Swing Adsorption (VSA) Vacuum swing adsorption (VSA), a modification of the PSA process, has been developed to improve the regeneration efficiency with lower power consumption and easier

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Figure 10.5 Vacuum swing adsorption (VSA) system.5

operating procedure.4,5,36e39 It differs from PSA by the use of low absolute pressures. This is completely dependent upon the nature of the gas component and its affinity to the sorbents. Figure 10.5 shows a typical apparatus for a vacuum swing adsorption system in which the inlet is slightly compressed (<1.5 atm). This process is composed of double adsorbentpacked reactors for adsorption and desorption, and a vacuum pump for recovering CO2. Figure 10.4b provides a three-step vacuum swing adsorption cycle as follows: (Step 1) adsorption of the inlet gas as it flows from bottom to top at fixed rate; (Step 2) evacuation, removing adsorbed components; and (Step 3) re-pressurization to atmospheric pressure with waste gas.5 The VSA cycle is quite similar to PSA, except for the absence of the pressurization step, which makes it simpler and probably more costeffective.

Chemical Adsorption Chemical adsorption, also known as chemisorption, on solid materials is achieved by substantial sharing of electrons between the surface of adsorbent and adsorbate to create a covalent or ionic bond. Thus, chemical adsorption may not be fully reversible, and could require high energy for regeneration. As carbon dioxide molecules are adsorbed on the surface of the adsorbent through valence bonds, they form a monolayer. The chemisorption energy between the gas molecules and adsorbents can vary significantly depending on the bond strength between adsorbent and adsorbate species. In terms of adsorption kinetics, adsorption rates usually involve activation energy barriers. If the activation energy is sufficiently high, practical equilibrium status may be obtained slowly.

CO2 Sorption

Figure 10.6 Process ﬂow diagram for CO2 capture by the temperature swing adsorption (TSA) system.40

Since solid sorbents are usually based on alkali and, or alkaline-earth metals, removal of adsorbed CO2 can be accomplished by calcination, as done in temperature swing adsorption (TSA).12,15,16,23,32,41e84 TSA is based on adsorption at low temperatures and desorption of adsorbed gas components at higher temperatures under constant pressure. Yamasaki et al. state that this adsorption method has the advantage of being environmentally friendly and requires low energy in the CO2 capture process.3 However, this is very dependent on the chemi-sorbent. A schematic diagram of a typical temperature swing adsorption system is outlined in Figure 10.6. The CO2-containing gas stream enters the reactor after passing the cooler and is adsorbed. The adsorbent is then heated, releasing CO2. The adsorbent is regenerated during desorption. During use, the chemical nature of the adsorbents can be modified by surface reaction or incomplete dissociations so that the original sorbents might not be fully regenerated after desorption. TSA is typically energy intensive due to the high energy consumed during the regeneration process. Additionally, TSA systems can have issues with impurity in the waste steam. For the optimum condition of the system, the high sorption affinity of the sorbent and improvement of processes should be considered, leading to cost-effective operation and maintenance.

ADSORBENT PERFORMANCE FOR CO2 CAPTURE Several different kinds of sorbents, e.g. amines, zeolites, potassium carbonate, lithium zirconate, sodium carbonate, and minerals, have been used for physical and chemical adsorption of CO2. The potential adsorbents must have: (1) high selectivity and adsorption capacity for carbon dioxide; (2) adequate adsorption/desorption kinetics;

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(3) stable cyclic adsorption capacity; (4) mechanical durability to maintain reasonable performance after extended cyclic exposure to high pressure or high temperature; and (5) low energy needs for regeneration of pure CO2.75

Physical Properties and Capture Capacities of Different Sorbents Physical properties that are the most important for characterizing solid sorbents include BET (Brunauer-Emmett-Teller) surface area and pore volume (Table 10.1). It can be seen that BET surface areas and pore volumes vary from 5 to 1,300 m2/g, and from 0.02 to 1.1 ml/g, respectively. Larger surface areas and larger pore volumes have long been regarded as key adsorbent properties. However, chemical properties are also relevant. Lithium zirconate, for instance, has a relatively low BET surface area (5 m2/g) and pore volume (0.02 ml/g), but its capture capacity is among the highest (0.29 g CO2/g sorbent). Table 10.2 shows the CO2 capture capacities of various solid sorbents. Typically, CO2 sorption capacities of conventional solid sorbents are of the order of ~0.1 g CO2/g sorbent. In order to reach the goal of ~0.15 g CO2/g sorbent, several studies have modified well-known sorbents in order to improve their capture capacities, which result from high surface areas, high CO2 affinity, and covalently tethered sorbents.12,53,54,85 Several sorbents are not effective in the presence of water vapor concentrations found in power plant flue gas, making them less attractive for use in such applications. Therefore, there is a need for materials with enhanced CO2 sorption capacities in the presence of gas streams containing water vapor. That would make them useful in realworld applications such as in coal-fired power plants.

Table 10.1 Physical properties of different solid sorbents Bulk density BET surface Solid sorbent (kg/m3) area (m2/g)

Tertiary amine (DBU) Aminated SBA-15 Zeolite 13X Zeolite/activated carbon Li2ZrO3 Activated carbon Hydrotalcites (HTlc) Mg-Al-CO3 HTlc Serpentine Metalorganic frameworks (MOFs)

689 500 860 440 2500 -

369 200~230 726 620 5 1300 271 184 330 345~2833

Pore volume (ml/g)

Reference

1.1 0.72 0.25 0.19 0.02 0.6e0.8 0.55 0.31 0.23 -

53 60 29 29 23 6,86 45 75 87 88

CO2 Sorption

Table 10.2 CO2 capture capacities of different solid sorbents Capacity (g CO2/ Solid Sorbent g sorbent) Gas composition 1

Aminated SBA-15

0.07~0.18

SBA-HA2

0.09

Tertiary amine

0.13

K-Li2ZrO3/ Li2ZrO3 MCM-413

0.22~0.29 0.09

K- HTlc4

0.44

Hydrotalcites (HTlc) Mg-Al-CO3 HTlc Zeolite 13X Zeolite/activated carbon Activated carbon Basic alumina Metalorganic frameworks (MOFs)

0.04 0.02 0.22 0.22 0.18 0.03 1.47~1.5

10% CO2, 88% He with ~2% H2O 15% CO2, 85% N2 with 20 mL/min rate of H2O 10% CO2, 88% He with ~2% H2O 10% CO2, 90% N2 15% CO2, 85% N2, 4% O2 15% CO2, 75% N2 with 10% H2O 11% CO2, 89% N2 20% CO2, 80% N2 15% CO2, 85% N2 15% CO2, 85% N2 17% CO2, 79% N2, 4% O2 100% CO2 -

References

60 53 12 23 89 18 45 75 29 29 6 44 85,88

1

SBA-15 denotes a type of mesoporous silica (Santa Barbara Amorphous type material, or SBA-15); SBA-HA Hyperbranched mesoporous aminosilica; 3 MCM-41, a type of mesoporous silica (Mobil Crystalline of Materials, or MCM-41); 4 K-HTlc potassium modified HTlc. 2

Zeolites In nature, zeolites, which are a class of regularly arrayed porous crystalline aluminosilicates of TO4 tetrahedra (T ¼ Si or Al), are formed by the reaction of volcanic rocks, ash, and water. Aluminum atoms in the supported silicate cause negative framework charges, leading to exchangeable cations such as alkali cations in the pore space. Due to the unique porous characteristics, such as molecular-size holes and pores of their framework functioning as molecular sieves, zeolites have high selectivity and capacity for CO2 adsorption. Not only can zeolites be artificially synthesized, but they are also inexpensive and environmentally friendly because of the natural existence of the materials and their use of physical sorption leading to an unchanged structure after sorption.5,13,20,22,29,37,41,47,52,59,64,67 These characteristics place zeolites among the top adsorbent materials. Different types of zeolites have been utilized in various adsorption methods such as PSA,6,13,19,20,22,25,27e31,41 VSA,5,36,37 and TSA52,59,62,64,67 with varying selectivity for CO2 over other gases such as CH4, CO, and N2. For instance, the selectivity for CO2 over N2 on zeolites in PSA is ~48% whereas in TSA and VSA it is ~90% and ~81%,

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respectively.30,59,90 Due to these high selectivities, the adsorption activity on zeolites can be affected by temperature and pressure; TSA and VSA on zeolite sorbents have been given more attention than when using PSA. However, PSA is still used to separate gas mixtures due to its cost-effectiveness and simple operation. Pure activity of sorbents for CO2 sorption leads to the modification of sorbent structure and the synthesis of different sorbents. Among the various types of zeolites, 13X and 5A have high CO2 capture capacities, due in part to their highly regular pore structure and large surface area.64 The kinetics of CO2 adsorption on zeolites reaches its equilibrium within a few minutes, being significantly faster in the initial adsorption period. Due to the presence of water vapor, CO2 adsorption in the presence of water vapor remains a relevant topic for the application of CO2 separation technologies to power plants. Li et al. studied CO2 capture capacities on zeolite 13X and the effect of water vapor in a VSA system at 30 C.5 The breakthrough curve of CO2 with water vapor showed a considerable effect of water vapor on the CO2 capture capacity (Figure 10.7), though the H2O adsorption capacity was hardly impacted by the presence of CO2.5,91 These results show a drawback in the use of this VSA system at 30 C in realistic flue gas mixtures since gas mixtures contain 8~10% water. To avoid water condensation and saturation over the adsorbent, the temperature could be held above 100 C, leading to its evaporation. However, that could damage the sorbent and cause pressure increases. Moreover, such high temperatures would increase operating costs. To evaluate the characteristics of zeolites, four important properties are discussed below.

Figure 10.7 Vacuum swing adsorption for CO2 þ air, H2O þ air and CO2/ H2O þ air in zeolite 13X at 30 C.5

CO2 Sorption

Zeolite Selectivity To capture CO2 effectively, one important property of the sorbent is high selectivity of CO2 over other gases, particularly for species with a smaller molecular volume, such as N2. In order to evaluate the sorbent selectivity for CO2 over N2, equilibrium adsorption isotherms of CO2 and N2 gases on the adsorbent are developed (Figure 10.8). The adsorption isotherms of CO2 and N2 on zeolite 13X were measured in the pressure range of 0e2.0 atm and the temperature range of 30e50 C.30 The amount of adsorbed CO2 on zeolite 13X was over four times higher than that of N2 in most of the temperature range, showing that zeolite 13X has a higher selectivity for CO2 than N2. Adsorption isotherms are modeled by the Langmuir and the Langmuir-Freundlich isotherms: Langmuir isotherm: qi ¼

qsi bi pi 1 þ bi pi

Eqn. 10.1

Figure 10.8 Evaluation of zeolite 13X on the basis of isotherm, purity and recovery in PSA. (a) Adsorption isotherm of CO2 (1) and N2 (2).29 (b) Effect of cycle time on the purity and recovery of N2.30

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Table 10.3 Langmuir Isotherm Parameters for CO2 and N2.27 Gas ai,1 (mol/g) ai,2 (K)

CO2 N2

1.95 1.03

2.03 0.61

bi,0 (1/mmHg)

bi,1(K)

87.84 2.81

2200 2520

Langmuir-Freundlich isotherm: 1=n

qi ¼

qsi bi pi

1=n

1 þ bi pi

Eqn. 10.2

where n is a power constant and the isotherm parameters bi and qsi are functions of temperature as follows: bi ¼ bio expðb1 =T Þ qsi ¼ ai;1 þ ai;2 =T

Eqn. 10.3

where bi denotes the Langmuir isotherm parameter (1/mmHg), qsi the saturation amount adsorbed of the ith component (mol/g), and bio, b1, ai,1 and ai,2 are calculated by isotherm experiments by using CO2 and N2 gases and zeolite 13X (Table 10.3). The Langmuir isotherm models the adsorption of gas molecules on the surface of a solid sorbent at constant temperature on a fixed absorption bed. It simulates the surface coverage of a monolayer without intermediate forces between gas molecules already adsorbed on the surface of sorbent. The kinetics parameters for CO2 and N2 can be found in Table 10.3. Purity and Recovery of Zeolite Another important property of zeolites is purity and recovery, calculated from data obtained in cyclic processes with the same sample. Purity is the ratio of the molar amount of a certain gas species produced over the total amount produced during the cyclic operation at steady state:30 R Tcycle v,yproduct dt Eqn. 10.4 Purityð%Þ ¼ 0 R Tcycle vdt 0 Tcycle ¼ ad/desorption cycle time, t ¼ reaction time, v ¼ feed rate, y ¼ the mole fraction of the component of interest. Recovery is the ratio of the molar amount of a given gas species in the product over the amount of the same species in the feed:30 R Tcycle vproduct ,yproduct dt Recoveryð%Þ ¼ 0 Eqn. 10.5 vfeed ,yfeed ,Tad

CO2 Sorption

Tcycle ¼ ad/desorption cycle time, t ¼ reaction time, v ¼ fluid velocity, y ¼ the mole fraction of the component of interest, Tad ¼ adsorption time. In a typical PSA processes, the CO2 purity increases with increase in cycle times, the operating time of ad/desorption is due to the accumulation of the molar amount of CO2. Thus, in order to operate the system efficiently, with enhanced CO2 purity and recovery, one needs to optimize various parameters such as operating cycle times, flow rate, adsorption time, product purge step time, and inert mole fraction (the ratio of inert gas species to total feed species).6,28,30 Based on the effect of cycle time on the purity and recovery of N2 (Figure 10.8b), the purity improved almost linearly with cycle time. An optimum cycle time was observed at ~300 seconds for zeolite 13X since the separation of CO2 and N2 on zeolites is equilibrium, rather than kinetically controlled, resulting from the favorable operation under relatively long ad/desorption times.30 The Cost of CO2 Separation Using PSA and VSA with Zeolite Capture cost is one of the most important factors to consider before using a process in realistic industrial applications. The capture cost of CO2 with a given adsorbent is given by:19 n P Ki þOi ð1þdÞi

i¼1 Cost of CO2 avoided ¼ P n ðCO2 i¼1

Eqn. 10.6

avoidedÞi ð1þdÞi

where Ki and Oi are the capital and operating costs (US$ million) in the ith year, d is the discount rate (% pa), and “CO2 avoided” is the annual amount of avoided CO2 emissions in million tons. CO2 avoided is defined as the difference between CO2 emitted and CO2 captured. The capital and operating costs include CO2 recovery and purity, and CO2 compression.19 The total capital cost is estimated as the cost of all process equipment in a given facility; the operating costs provide fixed general maintenance costs comprising labor, payable non-income government taxes, and insurance cost.19 Table 10.4 Estimated capacity cost and capital cost of CO2 capture with zeolite 13X. VSA19 PSA (IEA GHG)92 PSA19

Cost year CO2 recovery rate (%) CO2 purity (%) Energy penalty (%) Capture cost (US$/ton of CO2 avoided) Total capital cost (US$/kW)

2006 85 48 30 51

2006 85 48 35 57

1992 95 40 64

1.3

1.5

1.5

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Estimated CO2 capture cost for PSA and VSA with commercial zeolite 13X ranged from $51 to $57/ton of CO2 avoided (Table 10.4). Capital costs ranged from $1.3106 to $1.5106. The International Energy Agency Green House Gas (IEA GHG) committee studied CO2 adsorption costs with the PSA process applied to the flue gas of power plants.92 They estimated that, in 1992, capture cost and capital cost were US$64/ton of CO2 avoided and 1.5US$/MW, respectively. These results indicate that, for the cost reduction of adsorbent regeneration, VSA is potentially more cost-effective than PSA, since the reduced capture costs lead to the lower energy penalty from 40% to 35% or 30%. Zhang et al. studied purity, recovery, and capital cost for the VSA process with zeolite 13X in terms of initial CO2 concentration, temperature, and evacuation pressure (Figure 10.9). The purity and recovery of VSA increased as the inlet CO2 concentration increased (Figure 10.9a). Above 12% CO2, over 80% purity and 80% recovery were

Figure 10.9 CO2 adsorption on zeolite 13X in VSA.37 (a) Effect of initial concentration at 20C and evacuation pressure of 3 kPa. (b) Effect of temperature at 12.1% CO2 concentration and evacuation pressure of 3 kPa. (c) Isotherms for CO2 and N2 adsorption. (d) Effect of evacuation pressure at 15C and 12.1% CO2 concentration

CO2 Sorption

observed and the estimated process cost reduced from ~US$21/ton CO2 avoided to ~US$10/ton CO2 avoided. Regarding operating costs, the VSA process can be suitable for CO2 flue gas concentrations larger than 40% in water gas shift reactors in integrated gasification combined cycle (IGCC) systems. However, it might not be suitable for low CO2 streams (~3%) typical of gas turbine processes, because the capture cost of VSA increases dramatically as CO2 concentration decreases. To use this VSA process in IGCC systems, additional steps in VSA might be needed to optimize the design parameters, such as vacuum pressure intensity, temperature, and the development of appropriate adsorbents. Based on the effect of temperature on the VSA process, capture cost increased with temperature (Figure 10.9b). Purity also improved as the temperature increased, which is predicted by the adsorption isotherms, but the recovery trend was different from others (Figure 10.9b). It increased slightly, up to 35 C, and then gradually decreased due to the reduction of capture capacity of the zeolite 13X.37 In VSA, optimizing vacuum pressure is a key to cost-effective operation. The higher the evacuation pressure, the higher the power consumption and the lower recovery (Figure 10.9d). Adsorption isotherms for CO2 explained again the relationship between purity, recovery, and vacuum pressure (Figure 10.9c). At low pressure (0~6 kPa) in the whole temperature range, the increment levels in the isotherms showed steep slopes. This supports the fact that evacuation pressure at desired temperatures is the major factor for the selection of optimal purity and recovery in the VSA process. Equilibrium and Kinetic Adsorption Li et al. carried out adsorption equilibrium and kinetic separation for multi-gas stream (CO2, N2, O2 and CH4) on zeolite b in a TSA system (Figure 10.10). Between adsorption equilibrium and kinetic separation, equilibrium selectivity is mostly used for TSA processes. The adsorption equilibrium separation factor is a function of inverse temperature (Figure 10.10a). The adsorption equilibrium separation factor is defined as:47 ae;A=B ¼

xA =xB yA =yB

Eqn. 10.7

where xA, xB, yA and yB are the mole fractions of components A and B in adsorbed (x) and fluid phases (y) at equilibrium, respectively. For typical Langmuir isotherms, this factor is constant and is derived by the ratio of the Henry’s Law constants:47 aei;A=B ¼

KpA KpB

Eqn. 10.8

where KpA and KpB are the Henry’s Law coefficients of adsorption equilibrium for components A and B respectively.

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Figure 10.10 Adsorption equilibrium and kinetic separation factor for a TSA process with zeolite b.47 (a) Adsorption equilibrium separation factor as a function of inverse temperature. (b) Adsorption kinetic separation factor as a function of inverse temperature.

When the temperature increased, the equilibrium separation factor decreased, resulting from the relative magnitudes of the heats of adsorptions of the two gas species47. ðDH DH Þ

A B This phenomenon is captured by the van ’t Hoff equation where the value of RT 47 reduces as temperature, T, increases. Due to the higher Henry’s Law constant of CO2 among several gas components, the selectivity of CO2 over other gases dominates at higher temperatures. On the other hand, Figure 10.10b describes the adsorption kinetic separation factor as a function of inverse temperature. For the evaluation of diffusion processes, it is useful to use a kinetic separation factor, indicating the separation mechanism of any binary gas

CO2 Sorption

system. Similar to PSA systems, the kinetic separation factor is related to the diffusion process, which is the major mass transfer mechanism in TSA. The kinetic separation factor is defined as:47 ak;A=B ¼

Dc;A Dc;B

Eqn. 10.9

where Dc,A is the diffusivity within a crystal of component A (m2/s), Dc,B the diffusivity within a crystal of component B (m2/s), the diffusivity Dc given by the Arrhenius Ea equation: Dc ¼ D0 expðRT Þ with D0 being a pre-exponential factor (m2/s), Ea thediffusional activation energy (kJ/mol), R is the gas constant (0.08206 atm$L$mol1$K1), and T is the temperature in Kelvin. Kinetic separation factors for N2 over CO2 and O2 over CO2 decreased as temperature increased, whereas the separation factor of N2 over O2 increased slightly, depending on the relative magnitudes of the diffusion activation energies.47 These results indicate that at low temperature, separation processes of N2 over CO2 and O2 over CO2 can be kinetically dominant over zeolite b. However, there are a few issues with the use of zeolites. First, in the presence of water vapor in concentrations typical of flue gas from power plants, the CO2 capture capacity is smaller than that in the absence of water vapor, due to the higher affinity to zeolite surfaces for water molecules compared to CO2 molecules.43 Secondly, the CO2 sorption capacity of these sorbents decreases as temperature increases41,46,64,67 while many other sorbents show the opposite behavior. 4,12,54,55,69,75,93 Thus, in application to commercial systems, it is important to find the optimal temperature range.

Activated Carbon The meso- or microporous carbonaceous structures of activated carbons (AC) are widely considered as one of the strongest molecular sieves due to their highly adsorptive properties, which results from their potentially high surface area (over 1000 m2/g) and large pore volume (over 0.6 mL/g). In addition, AC has lower cost and higher hydrophobicity than zeolites. The adsorption isotherms of CO2 and N2 on activated carbon were developed in the pressure range of 0e2.0 atm and the temperature range of 15e55 C (Figure 10.11a). CO2 was adsorbed on activated carbon in amounts that were over four times larger than N2 over the entire temperature range, showing that activated carbon has a high selectivity for CO2 over N2. The Langmuir-Freundlich model (Eqn. 10.2) simulated the isotherms better than the ordinary Langmuir model (Eqn. 10.1) for both gases. This higher accuracy in describing the adsorption process is due to the larger number of parameters (three).94 The effect of different operating variables such as product purge and pressure equalization on the purity and recovery of CO2 in the PSA system is discussed in

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Figure 10.11 Evaluation of activated carbon on the basis of isotherm, purity and recovery.6 (a) Adsorption isotherm of CO2 (1) and N2 (2). (b) Purity and recovery of CO2 with various processes in the PSA system. (c) Effect of cycle time on the purity and recovery of CO2

(Figure 10.11b). For the typical PSA operating cycle, including pressurization, adsorption, depressurization, and regeneration, the purity and recovery of CO2 were ~50% and ~72%, respectively. The amount of N2 gas adsorbed during desorption in the depressurization step resulted in the differences in purity and recovery.6 PSA system with additional processes of product purge and pressure equalization produced similar performances for purity. When both product purge and pressure equalization were included, the purity of CO2 increased significantly to ~91%. However, the recovery decreased to ~46%, leading to overall poor performance because the purge step did not

CO2 Sorption

enhance the N2 desorption rate. Thus, the cost-effective process and high recovery of gas species for PSA uses only pressure equalization, as it yields similar performance for purity and recovery. With regard to the effect of ad/desorption cycle times on the purity and recovery of CO2 and N2, the purity increased with an increase in the cycle time, whereas the recovery decreased (Figure 10.11c).30 An optimum cycle time was observed at ~15 min. Recently, Park et al.20 investigated the PSA process using a double-layered sorbent bed with two adsorbents, activated carbon and zeolite 5A for multi-component (H2, CO2, CH4, and CO) adsorption. Based on the equilibrium isotherms of the four gas components on activated carbon (a) and zeolite 5A (b) (Figure 10.12), the selectivities of CO2 were significantly higher compared to the other gases on both adsorbents. However, strong CO2 adsorption at low pressures (high affinity) on zeolite 5A is an issue because a low pressure swing would not be effective. Since two adsorbents have a different affinity for inlet gas species in the double-layered column, an optimal design of layered column depending on feed quantities of each adsorbent is discussed. The short layer height of the activated carbon would lead to an unused region in the zeolite layer when CO2 began to break through the activated carbon; whereas the long layer height of the activated carbon would cause an unused region in the activated carbon layer when CO or CH4 broke through the entire column length.20 Therefore, the layer height of activated carbon should be optimized for the reduction of the unused column portion, and the feed amount for adsorption in the double-layered column should also be maximized at this optimal column height. When the activated

Figure 10.12 Equilibrium isotherms for a double-layer PSA system for multi-gas components.20 (a) Equilibrium isotherms on activated carbon at 299 K. (b) Equilibrium isotherms on zeolite 5A at 298 K.

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Figure 10.13 Effect of layer length on double layer PSA system for multi-gas components.20 Total bed length: 120 cm; QD: total inlet into layered bed; QAC: total inlet into a pure carbon bed, 14.8% vol CO2 at 17.0 atm.

carbon bed length was ~74 cm out of the total bed length of 120 cm, the rate of feed quantities on the double-layered sorbent reached a maximum (Figure 10.13). Thus, the optimal ratio of layered bed length for activated carbon over zeolite 5A was 0.617. According to the bidisperse pore models involving the micropore or macropore structures of activated carbons, the CO2 adsorption of activated carbons is dominated by diffusional mass transport. Because of their equilibrium capacity on the short kinetic time, the CO2 adsorption kinetics on activated carbons are comparable to those on zeolites. However, since activated carbons have a high affinity to water vapor, resulting in the low CO2 adsorption capacity, an additional modification of the surface functionalities of activated carbon is required for a feasible application, because the flue gas of coalfired power plants contains 8 to 12% water vapor.

Amine-modiﬁed Adsorbents Monoethanolamine (MEA) has been utilized to make detergents, emulsifiers, polishes, pharmaceuticals, corrosion inhibitors, and chemical intermediates.95 Moreover, the reaction of MEA with ammonia produces the chelating agent ethylenediamine (EDA).95 Both EDA and MEA have been utilized for CO2 separation by the gas and oil industries for several decades. Based on the thermodynamics of amine-modified supports, CO2 adsorption capacity decreases as the temperature increases due to the negative values of heat of adsorption and entropy change. For the evaluation of adsorption kinetics, heattransfer limitations are a concern for the adsorption reactivity of the different types of amines or inorganic-compound-supported amines. This is due to the high heat of

CO2 Sorption

CO2 adsorption on amines and the inhibited thermal exchange of many supports, leading to an increase in heat effects during adsorption kinetics. Gray et al. proposed a CO2 adsorption mechanism on amine or amine-modified adsorbents for primary, secondary, and tertiary alkanolamines, which react with dissolved CO2 to produce amine carbamates and carbonates.50 Solid amine CO2 sorbents provide similar reaction patterns with CO2, water vapor, and the amine functional group on its surface. DRIFTS and in situ infrared (IR) analysis support the proposed sorption pathways on aminated adsorbents. IR peaks correspond to carbamate, carbonate, and bicarbonate species produced via CO2 adsorption with amine functional groups. The mechanisms of the interaction between CO2 and amines depend on amine structure. The reaction mechanism for CO2 adsorption on primary, secondary, or hindered amines that contain an amine functional group surrounded by a crowded steric environment is as follows:96e98 CO2 þ AmH%AmHþ COO ðzwitterionÞ

R-10.1

AmHþ COO þ B/AmCOO þ BHþ ðcarbamateÞ

R-10.2

where AmH is an alkanolamine that contains hydroxyl (-OH) and amino (-NH2, -NHR, and -NR2, R ¼ C2H5) functional groups on an alkane backbone and B is a base. Primary and secondary amines can react with CO2 to form carbamate compounds through the production of zwitterionic intermediates. For the zwitterionic mechanism for the formation of carbamate suggested by Caplow,96 the lone pair of the amine functional group is attached to the carbon from CO2 to produce the zwitterion, and then a free base deprotonates the zwitterion to form the carbamate. In the aqueous phase, this base can be another amine, H2O, or OH. For secondary amine, the amino group is bound to a secondary or a tertiary carbon atom; and the mechanism of hindered amines is defined structurally as a primary amine such as 2-amino-2-methyl-1-propanol (CH3CCH3NH2CH2OH, AMP) in which the amino group is associated with a tertiary carbon atom, or a secondary amine such as 2-piperidineethanol (C7H15NO, PE), where the amino group is attached to a secondary or a tertiary carbon atom.99 Under dry conditions without H2O and OH, the maximum amine efficiency of an amine adsorbent is 0.5 mol CO2 per mol N, whereas under moist conditions with H2O acting as a base, the maximum amine efficiency is 1.0 mol CO2 per mol, indicating that an amine reacting with H2O or OH provides more capacity for CO2 capture. In the case of tertiary amines, the reaction mechanism for CO2 adsorption is different. Rather than a direct reaction with CO2, tertiary amines catalyze the production of bicarbonate. Donaldson and Nguyen100 and Vaidya Kenig98 et al. discuss the mechanism of the base-catalyzed hydration of CO2 for the reaction of CO2 with tertiary amines. The reaction mechanism for CO2 adsorption on tertiary amines is as follows:97,98,100

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AmH þ H2 O%AmHþ 2 þ OH

R-10.3

CO2 þ OH %HCO 3

R-10.4

þ AmHþ 2 þHCO3 %AmH2 ,HCO3

R-10.5

First, the tertiary amine dissociates H2O to produce a quaternary cationic species, and OH reacts with CO2 to produce the bicarbonate anion. Then, the ion of the protonated amine and bicarbonate are coupled. For example, the MEA equilibrium for CO2 capture in aqueous medium is: C2 H4 OHNH2 þ H2 O þ CO2 4C2 H4 OHNHþ 3 þ HCO3

(R-10.6)

Following this reaction pathway, one mole of CO2 reacts with one mole of MEA to form stable bicarbonate compounds. Forward rate constants for the reactions of CO2 with various amines, such as MEA, diethanolamine (DEA), triethanolamine (TEA), and methyldiethanolamine, (MDEA) range from 9 to 7600 mol/L$s (Table 10.5). Since the high rate of reaction results in the high rate of mass transfer, MEA is regarded as a strong potential sorbent for CO2 separation among the various amines, despite the relatively high energy requirements for CO2 desorption due to its high heat of reaction.51 CO2 sorption with MEA solutions has several drawbacks: (1) low CO2 capture capacity; (2) high energy consumption to maintain the appropriate temperature (up to 70% of the total operating costs in CO2 separation plants); (3) high equipment corrosion rate; and (4) amine degradation by SO2, NO2, HCl, HF, and O2 which leads to a high sorbent makeup rate.60,101e104 Therefore, attempts have been made with various types of solid amines, or solid amines modified by metal oxides, due to the weaknesses of the MEA-based solution process.4,12,16,32,50,56,60,77,101e103,105e112 Hiyoshi et al. investigated CO2 adsorption on amine modified SBA-15 in the presence of water vapor. 3-Aminopropyltriethoxysilane (denoted as APS), N-(2aminoethyl), 3-aminopropyltriethoxysilane (AEAPS), and (3-trimethoxysilylpropyl)diethylenetriamine (TA) were also utilized as grafting agents. The impregnated samples were denoted as APS/SBA (obtained by APS impregnation into SBA-15), AEAPS/SBA, and TA/SBA, respectively. The adsorption capacity of amine modified SBA-15 was much higher than that of APS/SBA (Figure 10.14a).43 Figure 10.14b Table 10.5 Typical values of forward rate constant for CO2eamine reactions.51 Amine type K1 (mols/L$s) Heat of reaction (kJ/mol)

MEA DEA TEA MDEA

7,600 1,500 16.8 9.2

72 65 62 53.2

CO2 Sorption

Figure 10.14 Adsorption of carbon dioxide on amine modiﬁed SBA-15.32,43 (a,b) With water vapor. (a) CO2 breakthrough proﬁle at 60 C over TA/SBA (B) and APS/SBA(i) (,): the mixture of 15% CO2, 12% H2O with N2 balance; total ﬂow rate: 30 cm3/min. TA is (3-trimethoxysilylpropyl)-diethylenetriamine, APS 3-aminopropyltriethoxysilane. (b) Relationship between amine efﬁciency and surface density of amine over 3-aminopropyltriethoxysilane (APS (C)), N-(2-aminoethyl)-3-aminopropyltriethoxysilane (AEAPS (:)),(3-trimethoxysilylpropyl)-diethylenetriamine (TA (A)). (c,d) Separation of carbon dioxide and methane on TEA grafted SBA-15. (c) CO2 Breakthrough proﬁle for SBA-15 without TEA modiﬁcation: mixture contained 9.74% CO2, 11.22% CH4 with He balance; total ﬂow rate: 100 cm3/min; pressure: 0.5 Mpa. 1: CH4; 2: CO2; 3: He. (d) CO2 Breakthrough proﬁle for SBA-15 with TEA-modiﬁcation: mixture contained 9.74% CO2, 11.22% CH4 with He balance; total ﬂow rate: 100 cm3/min; pressure: 0.5 Mpa. 1: CH4; 2: CO2; 3: He.

describes the relationship between amine efficiency and surface density of amine. Amine efficiency is defined by the quantity of absorbed CO2 (mmol/g) over amine content (mmol/g) at 60 C with CO2 (15%), H2O (12%) and balanced with N2 (74%). Surface density of amine is denoted as the number of nitrogen atom over 1 nm2 of the SBA-15 surface.

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The increase in the amine efficiency with different supports was accompanied by an increase in surface density of amine. The adsorption process on amine grafted SBA-15 forms ammonium carbonate as follows: 113 2R-NH2 þ CO2 4R-NHCOO þ R-NCOþ 3

R-10.7

As shown, two moles of amino groups react with one mole of CO2 to produce one mole of ammonium carbamate; isolated amine on the SBA-15 surface does not attract CO2 for adsorption, but densely tethered aminosilane provides an effective adsorption site.43 Amine efficiency, defined by the quantity of sorbed CO2 (mmol/g) over amine content (mmol/g) at 60 C with CO2 (15%), H2O (12%) and balanced with N2 (74%), was found to be nearly linear with N-atom surface density (the number of nitrogen atoms on the surface area) for different supports (Figure 10.14b), indicating that the length among tethered amines on the surface is an important parameter for CO2 adsorption. The increase in the surface density leads to an increase in the NeH bond and C¼O stretching intensities from ammonium carbonate. To evaluate the selectivity of another type of amine, TEA grafted SBA-15 was analyzed to find the breakthrough curves of CO2 and CH4 without TEA and with TEA (Figure 10.14c and d). Helium behaved quite differently than CO2 and CH4, but CO2 and CH4 were not quite suitable for separation in the case of SBA-15 without TEA modification due to the high affinities of carbon for amine. On the other hand, with TEA impregnation, the breakthrough time of methane became shorter, from ~100 s to ~40 s; and for carbon dioxide it became much longer, from ~230 s to ~775 s, indicating that TEA grafted SBA-15 enhances the separation capability for CO2 mixed with other gases.

Organic/Inorganic Hybrid Materials Recently, new organic/inorganic hybrid materials have shown high potential for carbon dioxide capture. This increases the accessible number of sorption sites per mass of sorbent, which allows smaller capture systems. Sorbents in this category include covalently amineimpregnated silica supports such as amine-modified silane and polymerization of reactive amine on silica. Due to the chemical nature of the structure, morphology, and spatial arrangement of individual species, a complex structural layer of organic materials can be easily changed by the grafted molecular structure of the components under slight modification of conditions. Inorganic supports can also easily become suitable supplementary materials to a complex structural layer of organic materials. Hicks et al. investigated the synthesis of covalently silica-grafted hyperbranched aminosilica (HAS) with a high loading of amine adsorption sites. Such organic/inorganic hybrid aminosilicas are easy to synthesize, have strong covalent organic/inorganic linkages, and are potentially low cost.53 HAS, which is created by polymerization of

CO2 Sorption

aziridine, is prepared by mixing the silica substrate with a precursor of the amine polymer in solution.53 The amine polymer is bound on the silica surface, forming a solid material that is filtered out and dried. The HAS synthesis produces SieOeC bonds through ringopening polymerization with surface OH groups on the silica surface. Aziridine can be polymerized off from silane-functionalized surfaces that contain nucleophilic reactive amines.114 Since HAS has a low surface area, mesoporous silica (SBA-15) supports are added to produce hybrid aminosilicas with a high-loading amine adsorption sites on a large surface area. This leads to high capture capacities because SBA-15 supports have strong covalent bonds between amines and supports, which leads to fully regenerable capacity during multi-cycles of ad/desorption.53 The common synthesis methods for covalently amine-impregnated adsorbents are the application of saline chemistry and polymerization. The schematic structures of various amines modified with SBA-15 result from physisorption, surface reaction, or surface polymerization (Figure 10.15).53 The grafting process of the hybrid aminosilica is achieved by accumulating alkoxy groups from a silane on the surface silanols to create covalent bonds to the silica surface through SieOeSi bonds, including organic functionalities through substituents on the silane.114 These silane-modified samples can be

Figure 10.15 Schematic structure of SBA-15 modiﬁed with various amines.53

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reformed to make large amine-loaded dendritic structures that are mainly tethered in the pore space.114 Organic polymers and inorganic materials have also been combined by mutual dispersion at the molecular level.115e118 Since sorbents with a three-dimensional structure of inorganic components, such as silica dispersed at the molecular dimension in organic components, provide more stability on dynamic intermediate species, the hybrid components can be readily regenerated for adsorption.118 Unlike many different kinds of amine-modified silica, organic/inorganic hybrid aminosilica with a high amine loading has good stability and recovery capacity for multi-cycles. However, the regeneration ability of the amine-tethered supports synthesized by the physisorption method is inhibited due to the physical tethering of low molecular weight of amines onto the supports. Which leads to a high loading of amine due to leaching of the organic polymers from the supports through weak linkages between amines and supports. If approaches can be found to make the organic polymers remain linked, the regeneration stability will improve.53 Other organic/inorganic hybrid materials that function as molecular sieves and sorbents which have received significant attention recently are the metaleorganic frameworks (MOFs) developed by Yaghi and coworkers.88 They are attractive because of their high thermal stability, their modifiable chemical functionality, and their crystalline, highly well-ordered porous structures.88 In addition, they can develop large numbers of combined networks of topologies, linkages of struts, and metaleligand bonds in the structure, resulting in the production of a variety of MOFs. MOFs typically provide three-dimensional organiceinorganic hybrid networks bound by different combinations of the metaleligand bonds, creating hubs and the linkage of struts consisting of organic compounds. MOF-5 (Zn4O(BDC)3(DMF)8(C6H5Cl), where BDC ¼ 1,4-benzenedicarboxylate and DMF ¼ N,N0 -dimethylformamide), which was first synthesized by Yaghi and coworkers, consisted of tetranuclear supertetrahedral clusters connected by bidentate BDC ligands into an octahedral arrangement.119 The DMF molecules were substituted with dichloromethane molecules, and the void sites were created by vacuum removing clusters. In terms of a cross section of framework schemes, MOFs provides many types such as square channels (MOF-2), pores decorated with open metal sites (MOF-505 and Cu3(BTC)2), and hexagonally packed cylindrical channels (MOF-74). In addition, interpenetration types of MOF are interpenetration (IRMOF-11), amino- and alkylfunctionalized pores (IRMOFs-3 and -6), and the extra-high-porosity frameworks IRMOF-1 and MOF-177 (Figure 10.16).88 Millward and Yaghi reported the adsorption equilibrium results at room temperature to evaluate the adsorption capacity for CO2 on different MOFs.88 For instance, MOF-177 has a higher adsorption capacity than conventional adsorbents, such as

CO2 Sorption

Figure 10.16 Different crystal structures of MOFs at room temperature. For each MOF, the framework formula, pore sizes, and surface area are shown.88

zeolite 13X and activated carbon. For the evaluation of MOFs, Figure 10.17 shows experimental adsorption isotherms and isotherms developed using a molecular model in IRMOF-1 (MOF-5). The molecular model can capture the reaction patterns and steps to account for the adsorption mechanisms for CO2 over a wide temperature range. Grand Canonical Monte Carlo (GCMC) simulation provides the modeling results. Adsorption isotherm models showed S-shaped isotherms, resulting from the very large pore sizes.85 Typically, many kinds of MOFs have the potential to adsorb CO2 in the range of 0.15 to 1.5 g CO2 per g of MOF.88 IRMOF-1 also has a high capacity of about 1.5 g CO2 per g. The molecular model shows good agreement with the experimental results. The rapid

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Figure 10.17 Comparison of GCMC simulations and experimental adsorption isotherms for CO2 in IRMOF-1.85

increase in the isotherm graph shown in the Figure 10.17, CO2 molecules mainly fill the pores of the networks of the MOF cavities with an increase in pressure. This profile trend, known as a type V isotherm, can lead to hysteresis, but the experimental and model results did not find this to be the case.120 The shape of the isotherm plots results from favorable electrostatic interactions between CO2 molecules.120

Lithium Zirconate Compared to alkaline-earth minerals such as CaO and MgO, alkali metal materials (e.g. Li2O, Na2O, K2O, Rb2O, and Cs2O) generally provide lower adsorption capacity and require more energy for carbonation and calcination. Another type of lithium-based oxide, however, lithium zirconate (Li2ZrO3), used for CO2 adsorption, has attracted attention as a promising sorbent at high temperatures (in the range 400e800 C) due to its highly stable adsorption reactivity and thermal stability. CO2 carbonation and calcination are expressed by the reversible equilibrium reaction: Li2 ZrO3ðsÞ þ CO2ðgÞ 4Li2 CO3ðsÞ þ ZrO2ðsÞ DH o ¼ 160 kJ=mol

R-10.8

Lithium zirconate has high capture capacity (~4.5 mmol/g) and good stability over many regeneration cycles.58,63,80 Moreover, it has high selectivity for CO2 and low affinity for N2.58 At atmospheric pressure, the forward reaction, carbonation, is thermodynamically favored up to 800 C. Calcination dominates at temperatures above 800 C. The XRD pattern for structure modification during sorption and desorption shows that CO2 carbonation using Li2ZrO3 is reversible.63 The CO2 adsorption rate increases as the CO2 diffusion rate increases during the formation of eutectic molten carbonate in the range 400e800 C.63,80 Fauth et al. examined Li2ZrO3 carbonation modified with alkali (Li, K, and Na) as well as alkali/alkaline (K, Mg) carbonate.57 Binary

CO2 Sorption

Figure 10.18 Schematic diagram of the proposed mechanism for the CO2 adsorption on Li2ZrO3.63 (a) and (b) are the adsorption process on pure Li2ZrO3, and (c) and (d) on modiﬁed Li2ZrO3.

K2CO3/MgCO3 and KF with pure Li2ZrO3 had a faster CO2 sorption rate at 500 C than pure Li2ZrO3, the additional carbonation reaction with potassium due to Li2ZrO3, modified by NaF/K2CO3 with Na2CO3, had the highest capture capacity and the fastest CO2 sorption rate at 600 and 700 C. Ida and Lin described the proposed double-shell mechanism for CO2 adsorption and desorption on Li2ZrO3 (Figures 10.18 and 10.19, respectively). For adsorption, the relatively fast CO2 sorption may initially occur before the formation of a dense Li2CO3 shell on the surface and a dense solid ZrO2 shell in the core surrounded by the unreacted solid Li2ZrO3.63 After the formation of the Li2CO3 and ZrO2 shells, the carbonates

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Figure 10.19 Schematic diagram of the proposed mechanism for the CO2 desorption on Li2ZrO3.63

thicken, inhibiting CO2 sorption into the unreacted pore regions of materials. Thus, the process rate decreases although the sample continues to react in the pores of the sorbents. Additional reactions, on the surface of the unreacted Li2ZrO3 produces Liþ and O2 ions, and CO2 may diffuse interstitially through the Li2CO3 shell and react with Liþ and O2 on the surface of the ZrO2 shell. O2 can fill up the open sites and Liþ can diffuse through the ZrO2 shell since ZrO2 provides a large number of oxygen vacancies in its crystals.63,121 The ion sizes of Liþ and O2 are considerably smaller than those of CO2, resulting in a diffusion rate of CO2 into the Li2CO3 layer that is slower than that of both ions; thus, CO2 diffusion is the rate-limiting step.

CO2 Sorption

For modified Li2ZrO3 (Figure 10.18c,d), the Li2CO3/K2CO3 layer encompasses Li2ZrO3 at high temperatures, and changes into molten carbonates due to its lower melting point. The CO2 diffuses through this layer and reacts with Li2ZrO3, since CO2 diffusion in the molten carbonate is much faster than in the solid carbonate.63,122 At the same time, Li2CO3, which is immediately produced, combines with the molten carbonate. This results in an increase in the volume of the molten carbonate layer, which can lead to a higher CO2 sorption rate for modified Li2ZrO3. Therefore, the carbonation reaction is limited by Li2CO3 formation as well as the ion transport in the ZrO2 shell. For the CO2 desorption process (Figure 10.19), the experimental temperature increases to 780 C, which is higher than the melting point of Li2CO3; thus both Li2CO3 in the Li2ZrO3 and the lithium/potassium carbonate mixture in the modified Li2ZrO3 are formed in the molten phase (Figure 10.19b).63 First, Li2CO3 reacts with ZrO2 to form Li2ZrO3 and CO2 reversibly. Second, CO2 diffuses through liquid Li2CO3 or Li2CO3/K2CO3 mixture to the outside, and during the formation of the dense Li2ZrO3 shell, Liþ and O2 keep diffusing through the Li2ZrO3 shell for CO2 desorption.63 Consequently, Ida, Lin, and Fauth et al. found that the CO2 sorption rate depends on CO2 diffusion and Liþ and O2- ion transport within the outer layer of the solid Li2O3 shell.57,63 Further improvement of the CO2 adsorption rate is needed because the adsorption kinetics are too slow to make practical systems based on such sorbents. In addition the high temperature required for the carbonation reaction leads to a high energy requirement. The reaction rate of these materials is slower than that of sorbents for physisorption such as zeolites or activated carbons. Lin and co-workers found that CO2 adsorption with pure Li2ZrO3 at 673 K does not equilibrate after 13 days of reaction and the adsorption halftime was 12 hours.80 To help increase the CO2 sorption rate, dopants such as Li2O3/K2CO3 and Li2O3/Na2CO3 have been used, resulting in significant improvements. Additional treatments leading to higher surface area, greater pore volume, and smaller particle size can enhance the adsorption rate.

Mineral-based Adsorbents Magnesium-based and calcium-based minerals have been considered as sorbents for CO2 capture from gas streams.123e125 Mineral carbonation provides enormous storage capacities and is potentially environmentally friendly, and the sorbent reactions are naturally exothermal. Moreover, solid by-products can be beneficially used for mine reclamation or soil amendments, because magnesite (MgCO3) and silica (SiO2) exist naturally as serpentinized rocks.87 These potentially useful by-products can contribute significantly to improve cost-effectiveness.87,125 They can also enable permanent, sequestration and avoid the costs of transporting and pumping CO2 underground. Calcium and magnesium are not naturally available as oxides, but are typically found as silicate or aluminosilicate. Large amounts of magnesium-based silicate

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minerals, such as olivine ((MgFe)2SiO4) and serpentine (Mg3Si2O5(OH)4), are widely distributed around the world, particularly in coastal areas.123 Over geological time scales, olivine is eventually transformed into serpentine and magnesite (MgCO3), suggesting that serpentine is more ubiquitous than olivine. New CO2 capture by mineral-based solids are due to the reaction of magnesium- or calcium-rich minerals with CO2 to produce stable carbonates. The general reaction of olivine and serpentine minerals with CO2 is:87 ðMg; CaÞx Siy Oxþ2y þ xCO2 /xðMg; CaÞCO3 þ ySiO2

R-10.9

Since the mineral carbonation reactions of both olivine and serpentine are exothermic, the heat released can be used.123,124 Magnesite (MgCO3) and silica (SiO2) exist naturally as serpentinized rocks. During the carbonation reaction, filling and pluggage of the pores are observed and the carbonate layer thickens, prohibiting CO2 sorption into the unreacted pore sites of materials. Thus, the process rate is decreased, although the sample continues to react at the pores of the sorbents. After the fast carbonation period, structural modification of the sorbent could occur during the formation of products with subsequent CO2 capture. Even though they are thermodynamically favorable, the mineral carbonation reactions are slow near ambient temperatures around 25 C.48,87,124 Therefore, a significant speeding of the carbonation reaction rate is needed for the development of an efficient system that would allow the feasible application of these sorbents. In order to improve the efficiency of CO2 capture of these minerals, several pretreatment processes such as magnetic separation, heat treatment, and acid treatment have been developed. Faster reaction rates can be achieved by first dissolving the minerals in HCl solution. First, serpentine easily dissolves in an HCl solution. Second, water is added to precipitate solid Mg(OH)2, resulting from the reaction (R-10.14). However, this process is too energy intensive and requires large quantities of water. To reduce the water requirement, the serpentine can be dissolved in molten MgCl2 salts. CO2 can be added to this solution to form MgCO3 and the molten MgCl2 salts can be recycled.48 However, this method uses highly corrosive chemicals, and operation and maintenance difficulties ensure. Current research has focused on the dissolution of the serpentine without using acid. One direction is to add sodium carbonate (Na2CO3) to the solution. The process is pressurized to 30 atm, and MgCO3 is formed via precipitation. However, dewatering is required, leading to energy penalties of the order of 20% of the electricity output of a power plant.48 Thermodynamics and Kinetics Magnesium and calcium silicates react with CO2 to form calcium and magnesium carbonates.

CO2 Sorption

These reactions are naturally exothermic as follows: Mg3 Si2 O5 ðOHÞ4 ðsÞ þ 3CO2 ðgÞ/3MgCO3 ðsÞ þ 2SiO2 ðsÞ þ 2H2 OðlÞ ðDHr ¼ 64 kJ=molÞ

R-10.10

Mg2 SiO4 ðsÞ þ 2CO2 ðgÞ/2MgCO3 ðsÞ þ SiO2 ðsÞ ðDHr ¼ 90 kJ=molÞ R-10.11 While the reactions are thermodynamically favorable, the reaction kinetics are slow. This is the most critical challenge. The slow natural carbonation kinetics can be enhanced by increasing the surface area, reaction temperature, and pressure, or by using a catalyst. Surface carbonation, which occurs in the direct solidegas phase reaction, is relatively uncomplicated. Solid calcium and magnesium minerals, including oxides or hydroxides, produce surface-adsorbed carbonates without any liquid phase. Compared to MgO and Mg(OH)2, the carbonation of CaO and Ca(OH)2 proceeds relatively fast in a forward reaction (at high enough pH), occurring within several minutes at a high temperature (< 680 K) and pressure (at 1 bar).126 Despite the rapid reaction, the feasible application of calcium-based sorbents as feedstock is limited due to their relatively low quantity as naturally occurring ores. Mg(OH)2 dehydroxylation to MgO has been broadly used due to its importance in industrial MgO production, resulting from its simple chemical property and structure. In particular, this process leads to MgO carbonation to capture CO2. The reaction mechanism of the carbonation of Mg(OH)2 is: MgðOHÞ2 ðsÞ/MgOðsÞ þ H2 OðgÞ ðdehydroxylationÞ MgOðsÞ þ CO2 ðgÞ/MgCO3 ðsÞ ðcarbonationÞ

R-10.12 R-10.13

The first step (R-10.12) is reversible: MgOðsÞ þ H2 OðgÞ/MgðOHÞ2 ðsÞ ðrehydroxylationÞ

R-10.14

Final carbonation proceeds through the intermediate formation of magnesium oxide layers. During the surface carbonation process, a thin layer of the carbonates forms as a diffusion barrier to both the outward diffusion of H2O and the inward diffusion of CO2, which decreases the process rate. Kinetic Improvements Given the nature of the slow reaction kinetics, additional process, and sorbent developments are required for commercial application. Balaz et al. investigated the enhancement of the chemical reaction rate through mechanical activation, which can potentially affect heterogeneous reactions by the formation of defects, increasing surface

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Table 10.6 The milling parameters and physicochemical properties of olivine samples.74 Speciﬁc Revolutions Milling time H2O addition surface Milling energy E Type of mill (revs/min) (min) (ml) area (m2/g) (kWh/kg)

Raw Attritor Planetary mill Nutating mill

1,500 450 900

120 30 10

100 500

0.25 35.2 5.2 9.4

2.31 6.03 0.51

area, and delocalization.74 Mechanical activation of olivine samples was performed in the three high-energy mills: (1) a laboratory attritor mill was run at rotation (1,500 revs/min) for milling times of 5e120 min and was operated at ambient temperature and free access of air; (2) a laboratory planetary mill was run at the planet carrier rotation speed (450 revs/min) for milling times of 2e30 min and also operated at ambient temperature and free access of air; (3) an industrial nutating mill was run at 900 revs/min, and operated in a dry mode, at ambient temperature and free access of air. In several tests, the milling was carried out with the presence of water (50, 100, 200 and 500 ml).74 During the high-energy milling processes, particles are crushed, rounded, and made smaller, leading to larger surface areas (Table 10.6). Attritor milling was the most promising, yielding the largest improvement in olivine surface area (from 0.25 m2/g to 35.2 m2/g). As the application of nanomaterials extends to various areas of science and engineering, it has also been applied to carbon dioxide capture. Wu et al. carried out CO2 adsorption at 600 C and desorption at 800 C on nano-CaO/Al2O3 adsorbent (Figure 10.20). The adsorption ratio, defined as the percentage of adsorbed CO2 within 10 min, began at 100%, then stabilized at about 75% after 15, cycles.69 During cycling, the average pore size increased ~75.9%, whereas pore volume and surface area significantly decreased (~56.7% and ~73.5%, respectively) due to the sintering of nanoparticles, and the formation of new products with larger particle size (Table 10.7). The desorption rate of CO2 from nano-CaCO3/Al2O3 was higher than that of micro-CaCO3/Al2O3 due to the surface energy of adsorbents improved with the decrease in adsorbent particle sizes. High surface energy at reaction sites leads to a high heat of hydration. This suggests that nano-CaCO3/Al2O3 can be a suitable CO2 adsorbent.

Sodium-based Sorbents Since water vapor is present in flue gas from power plants, alkali metal carbonates, such as K2CO3 and Na2CO3, have received some attention since they easily react with carbon dioxide and water vapor to form alkali metal hydrogen carbonate. In addition, the reactions can be reversed at relatively low temperatures in order to regenerate the sorbent.

CO2 Sorption

Figure 10.20 Adsorption and desorption on nano CaO/Al2Co3 sorbent.69 (a) Adsorption ratio of carbonation and regeneration for 50 cycles. (b) Comparison of desorption proﬁle dependent on temperature on nano- and micro-CaCO3/Al2Co3 sorbents.

Seo et al.55 carried out the CO2 carbonation and calcination with water vapor on a sodium-based solid sorbent. The reaction sequence is: Na2 CO3 ðsÞ þ CO2 ðgÞ þ H2 OðgÞ%2NaHCO3 ðsÞ DH 0 ¼ 135 kJ=mol R-10.15 Since the heat release could reverse the carbonation process during the highly exothermic chemical adsorption, this study used a fluidized-bed reactor providing high mass and heat transfer rate. It was conducted at 50 C using a simulated flue gas mixture (10% CO2, 77.8% N2 and 12.2% H2O) and the regeneration was carried out at 135 C using pure N2. Sorb NX35, consisting of 35% Na2CO3, and 65% support and binder, Table 10.7 BET analyses of CaO/Al2CO3 sorbent before and after 20 cyclic carbonations.69 Pore size (nm) Total pore volume (mL/g) Surface area (m2/g)

Before After

8.631 15.18

0.2578 0.1116

115.5 30.57

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was used with water vapor pretreatment in the bubbling fluidized-bed reactor using the TSA system (Figure 10.21a). It was suggested that using a Sorb NX35 adsorbent with water pretreatment has a potential for CO2 separation due to a high capture capacity of 8.1 g CO2/g adsorbent and 90% regeneration capacity. Water vapor pretreatment affects the capture capacity of this material (Figure 10.21b). Without water vapor pretreatment, CO2 reduction decreased exponentially; while with water vapor pretreatment, capture increased with the duration of

Figure 10.21 Temperature swing adsorption (TSA) system using sodium-based solid sorbent.55 (a) CO2 breakthrough proﬁle. (b) Effect of water vapor pretreatment. (c) Cyclic adsorption proﬁle. (d) Properties changes of adsorbent during cycle. Sample 1: dried fresh sample at 300 C; sample 2: 50 C carbonation for 1 h after 30 min pretreatment; sample 3: 135 C regeneration for 1 h under N2 after sample 2; sample 4: 300 C regenerated for 1 h under N2 after sample 2.

CO2 Sorption

the pretreatment. This result suggests that water vapor could play a catalytic role in the reaction of CO2 with Sorb NX35, instead of being just a reactant in the reaction (Na2CO3 þCO2 þH2O 4 2NaHCO3). However, overexposure to water pretreatment above the appropriate duration resulted in the loss of CO2 capture capacity and regeneration level. This is due to damp agglomeration of the adsorbent and a higher affinity of the water molecules to the sorbent surface compared to carbon dioxide molecules. During cyclic adsorption, CO2 removal capacities decayed sharply when regeneration was carried out at 135 C (Figure 10.21c). This could be due to water vapor accumulation inside the sorbent pores, resulting in poor contact between the CO2 molecules and the sorbent. When regeneration was carried out at 300 C instead of 135 C, the CO2 removal performance became close to the initial cycle, due to water removal from the sorbent and more favorable temperature. Since the properties of Sorb NX35 change during the carbonationeregeneration cycle, the BET surface area and pore volume were evaluated before and after reaction (Figure 10.21d). After adsorption, surface area and pore volume reduced significantly. These properties increased after desorption with a pure N2 stream at 135 C, although it is still lower than with the fresh sample. Finally, after regeneration at a higher temperature (300 C), they were fully recovered.

Hydrotalcite-like Compounds (HTlcs) Hydrotalcite-like compounds (HTlcs) are anionic and basic clays, known as layered double hydroxides (LDHs). These consist of positively charged brucite (magnesium hydroxide)like layers with negatively charged interlayers compensating anions and water molecules, resulting in a charge-balanced framework.75,127 They have been utilized as catalysts, adsorbents, catalyst supports, ion exchangers, filters, decolorizing agents, and stabilizers.45 In recent years, studies have dealt with application of these materials for CO2 adsorption because of their stable anion exchange, highly stable adsorption behavior, high mobility of both anions and water molecules, resulting from large surface areas, and the stability of the frame structure.45 The general formula of HTlcs is represented xþ nx2þ 2þ and M3þ are divalent as follows: ½M2þ 1-x Mx ðOHÞ2 ,½Ax=n ,mH2 O , where M 2þ 2þ 2þ 2þ cations, Mg , Zn , Ni , or Cu , and trivalent cations, Al3þ, Fe3þ, or Cr3þ, 2 respectively; An are interlayer anions, CO2 3 , SO4 , NO3 , Cl or OH , and the x value is typically (but not necessarily) in the range 0.17~0.33. HTlcs consist of positively charged bidimensional layers of metal oxides (e.g. MgO) or hydroxides (e.g. Mg(OH)2) with trivalent and divalent cations at the center of the octahedral sites in the hydroxide layer and anions in the interlayer (Figure 10.22a). The excess positive charges obtained by the replacement of trivalent cations are compensated by anions and water molecules in the interlayer zone, leading to a charge-balanced structure.45 To enhance the adsorption activity, the chemical structures of hydrotalcites

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Figure 10.22 Carbon dioxide adsorption on hydrotalcite-like compounds (HTlcs).45,75 (a) Schematic structure of hydrotalcite-like compounds. (b) CO2 adsorption dependent on the anion type at 300 C.

are modified by interchanging the metal composition of the framework, or combining dopants such as alkali metals, preventing sintering of the particle. CO2 adsorption capacities on HTlc sorbents significantly depend on the type and amount of divalent cations, trivalent cations, and anions. Based on the carbon dioxide and OH) adsorption for different samples with two types of anions (CO2 3 2 (Figure 10.22b), HTlcs with CO3 have higher capture capacities than those with OH. HTlcs with the carbonate ion have more void space and larger interlayer spacing than the ones with OH (0.765 mm vs. 0.755 mm), where they can adsorb more carbon dioxide molecules.45,128 Yong et al. proposed that the optimized composition ratio of divalent (Mg2þ) and trivalent (Al3þ) cations was a 1:1 ratio, producing a maximum capacity of CO2 adsorption. Incorporated aluminum enhances layer charges in HTlcs; hence, the increase in the amount of aluminum results in the reduction of the interlayer spacing, leading to a decrease in the adsorption sites. When the type of anion was altered from OH to

CO2 Sorption

Table 10.8 CO2 adsorption capacities on Mg-Al-CO3 layered double oxide (LDO) dependent on temperatures at 100 kPa.93 Combined Reversible Irreversible Reversible Temperature (mmol/g) (mmol/g) (mmol/g) (%) ( C)

100 200 300 400

0.231 0.486 0.249 0.169

0.192 0.429 0.208 0.132

0.039 0.057 0.041 0.037

87.0 88.3 83.5 78.1

CO2 3 , adsorption capacities increased from 0.25 to 0.5 mmol/g at 300 C and 1 bar due to larger interlayer spacing and higher layer charge in the HTlcs with CO2 3 anions. Reddy et al. showed combined adsorption (both physisorption and chemisorption), as well as reversible (physical), and irreversible (chemical) CO2 adsorptions on Mg-AlCO3 layered double oxide (LDO) (Table 10.8).93 Adsorption capacities typically decreased with increasing temperatures. However, the highest capture capacity was observed at 200 C instead of 100 C, because physical adsorption dominates at 100 C. This indicates that CO2 adsorption on Mg-Al-CO3 layered double oxide (LDO) results from both physical and chemical adsorption on different sites.75 Over the temperature range studied, reversible adsorption constituted over 78% of the overall adsorption. At higher temperatures, or higher kinetic energies, desorption from the sorbent surface becomes favorable.

CO2 CAPTURE DEVELOPMENT STRATEGIES CO2 capture by the reduction of anthropogenic CO2 emissions from large point sources has gained significant attention as a way to reduce greenhouse gas emissions throughout the world. Among various separation technologies, including absorption, adsorption, membrane separation, cryogenic distillation, and chemical-looping; adsorption using solid sorbents is one of the most promising. Adsorption, a heterogeneous process, is one of the methods that can separate CO2 from the flue gases of coal-fired power plants. Due to interaction between sorbent and CO2 gas species, CO2 molecules are captured on the surface of the sorbents. In general, CO2 adsorption processes with various adsorbents show highly stable and cyclic capture capacities and low energy consumption for regeneration in comparison with aqueous systems. Adsorption kinetics depend on temperature, pressure, interaction energy between sorbent and CO2 on the surface, and pore size or surface area of the adsorbents, corroborating the fact that many parameters affect capture capacity, operation cost, and process efficiency. Some of the issues with the adsorption processes are, limited CO2 selectivity to separate CO2 from realistic flue gases, and low adsorption rates. In order to overcome these issues, current research efforts are focusing on developing solid sorbents with fast

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ad/desorption kinetics, high selectivity, high capture capacity, high recycle stability, and targeting realistic applications. Further adsorbent research is needed to, understand the mechanism of CO2 adsorption on the various solid adsorbents; improve the kinetics of adsorption and desorption of CO2 over different sorbents; and modify sorbents to optimize their performance.42 In addition, the adsorption technologies can be improved by process modifications such as: (1) combined temperature and pressure swing adsorption (TPSA), pressure swing adsorption using a temperature swing system to improve recovery; (2) fractionated vacuum-pressure swing adsorption (FVPSA), a dual-bed process which simultaneously produces a 98þ% nitrogen-enriched gas and an 80e90% oxygenenriched gas from ambient air to improve the CO2 purity as well as the N2 purity;13 (3) combined temperature and vacuum swing adsorption (TVSA); (4) insertion of more steps in the PSA and VSA processes; and (5) addition of multi-columns for the adsorption and desorption process. Finally, cost-effective scalabilty to megaton CO2 capture fluxes is critical to enable CCS and transition it from bench-top to a real power plant scale (e.g. the German power plant pilot study). In summary, the progress in the adsorption technology which has been reviewed can make it a competitive approach for CO2 separation from gas streams based on both economic aspects and technical performance.

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11

Ammonia and Derivatives Contents Historical Background Feedstocks for Ammonia Synthesis Ammonia Synthesis Equilibria and the Synthesis Loop Ammonia Synthesis Catalysts Ammonia Synthesis Kinetics Ammonia as Nitrogen Fertilizer Ammonia as a Transportation Fuel Urea Nitric Acid and Ammonium Nitrate Future of Nitrogen Fertilizers References

341 342 342 344 345 347 348 348 349 350 351

HISTORICAL BACKGROUND Nitrogen is one of the three primary plant nutrients. A bag of fertilizer typically lists three numbers. The first is the percent nitrogen, the second the percent phosphorus (P2O5 equivalent) and the third is the percent potassium (K2O equivalent). Nitrogen is abundant in the natural world, but much of it is in an unreactive form, N2, in the atmosphere. In unfertilized soils, nitrogen is often a limiting nutrient. A breakthrough occurred when Fritz Haber, a German chemist, discovered that he could synthesize ammonia directly from N2 and H2 using transition metal catalysts at high pressure. N2 ðgÞ þ 3H2 ðgÞ/2NH3 ðgÞ

DH rxn ¼ 91:9 kJ=gmole

R-11.1

Haber won the 1918 Nobel Prize in Chemistry for this discovery. Carl Bosch converted Haber’s discovery into an industrial process, so the ammonia synthesis process is often known as the Haber-Bosch process. The first commercial plant was commissioned in 1913, using an osmium catalyst.1 The Haber-Bosch process had the unintended consequence of prolonging World War I. Earlier, nitrogen compounds for explosives came primarily from sodium nitrate mined in Chile. During World I, Germany was cut off from the Chilean mines, and turned to synthetic ammonia for the production of explosives. Today, nitrogen fertilizers such as ammonia, ammonium nitrate, ammonium phosphates, and urea constitute the bulk of synthetic nitrogen compounds. All of these compounds, plus nitrogen-based explosives, polymers, and other nitrogen-containing synthetic chemicals, start with the synthesis of ammonia. Coal Gasification and Its Applications. ISBN B978-0-8155-2049-8.10011-7, doi:10.1016/B978-0-8155-2049-8.10011-7

Ó 2011 Elsevier Inc. All rights reserved.

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FEEDSTOCKS FOR AMMONIA SYNTHESIS Chapter 7 described the production of H2 for ammonia synthesis from syngas. The air separation unit (ASU) used to generate oxygen for the gasification process, produces more than enough N2 for ammonia synthesis. Ammonia plants were initially designed to produce ammonia from coal. When natural gas became abundant, natural gas largely replaced coal as a feedstock. In the last several years, ammonia production has shifted to countries with large natural gas production and relatively small local natural gas consumption. Of course, natural gas can be liquefied and shipped in transoceanic tankers, but the higher normal boiling point of ammonia, 195 K (78 oC), versus methane, 111 K (162 oC), means that it is easier to liquefy and ship ammonia than natural gas. Rising natural gas prices could prompt a shift from natural gas feedstock back to coal. Hydrogen is produced from natural gas using the steam methane reforming reaction: CH4 ðgÞ þ H2 OðgÞ4COðgÞ þ 3H2 ðgÞ

DH rxn ¼ 206:2 kJ=gmole

R-11.2

Steam and natural gas are fed to tubes packed with catalyst in a furnace.2 External combustion, usually a natural gas-fed flame, provides the endothermic heat of reaction. The effluent of this reactor, which contains unreacted natural gas, is then fed to a secondary reformer along with air. Parital combustion of the gas provides additional heat for the endothermic steam methane reforming reaction. The air also provides nitrogen for ammonia synthesis. The air feed to the secondary reformer is adjusted so that the processed syngas contains the 1:3 N2/H2 ratio needed for ammonia synthesis. A similar approach could be used with coal gasification. A combination of air and O2 from an ASU could be adjusted so that the processed syngas has a 1:3 N2/H2 ratio. This would reduce the size of the ASU needed for the plant, as some of the oxygen would be provided by the air feed.

AMMONIA SYNTHESIS EQUILIBRIA AND THE SYNTHESIS LOOP Using the stoichiometry shown in R-11.1, an equilibrium constant can be written as: KNH3 ¼

ðpNH3 Þ2 ðpH2 Þ3 pN2

Eqn. 11.1

Calculated values for this equilibrium constant between 350 and 500 oC and 10 and 30 MPa are shown in Figure 11.1. Since four moles of gas react to produce two moles of gas, higher pressures push the reaction to the right. Higher temperatures reduce the equilibrium constant, which is typical for an exothermic reaction. Higher pressures favor greater conversion of the reactants to ammonia, but compression costs, and the cost of high pressure vessels limit the operating pressure.

Ammonia and Derivatives

Figure 11.1 Equilibrium constant for the ammonia synthesis reaction , as deﬁned in Eqn. 11.1, versus temperature and pressure. The equilibrium constant units are MPa2. Equilibrium constants were calculated using a Peng-Robinson equation of state and assuming a pure, 1:3 H2/N2 feedstock.

Lower temperatures also favor the equilibrium, but higher temperatures increase the rate of reaction; so the operating temperature is a compromise between favorable equilibria and reaction rates. Jennings and Ward1 show a typical operating pressure of 22 MPa. Gas enters the reactor at 400 oC, and is cooled between beds to limit the temperature rise due to the exothermic heat of reaction. Gas leaves the reactor at 470 oC. At the reactor exit temperature and pressure, with a pure, stoichiometric ratio feed stock, the equilibrium gas composition is 56.4% H2, 18.8% N2, and 24.8% NH3. This corresponds to 40% single-pass conversion of the feedstock. Since less than half the feedstock is converted to ammonia per pass. Ammonia is separated from the reactor effluent and the unreacted feed gas is recycled. This leads to the flow sheet shown in Figure 11.2. Feed gas is compressed, mixed with recycle gas, and preheated by exchanging the feed gas with the hot reactor effluent gas. The reactor shown in Figure 11.2 is a quench design. The bulk of the feed gas, at about 400 oC, is injected into the top of the reactor. Relatively cool quench gas, at about 150 oC, is injected between the catalyst beds to drop the temperature of the reactant gases. Other reactor designs are used, and all of these reactors are designed to remove heat to maintain the reaction temperature within a desired range. The effluent gas is cooled by a series of heat exchangers to condense ammonia. The uncondensed gas is recompressed and fed back to the reactor.

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synthesis reactor

feed gas 400 °C

quench gas 150 °C

cooling water

refrigerant

470 °C

- 5 °C liquid ammonia

steam

recycle gas

fresh feed gas

feed compresor

recycle compressor

Figure 11.2 Ammonia synthesis loop, showing a quench reactor.

Inert gasses accumulate in the recycle gas. One of these gases is argon, which is present in air at about 0.9%, and is a common impurity in O2 from ASUs. As discussed in Chapter 7, one of the methods used to eliminate CO and CO2, which are catalyst poisons, is to convert them to CH4, which behaves as an inert in the ammonia synthesis loop. A syngas purge is required to prevent excessive build up of inert gases. A hydrogen permeable membrane is typically used to recover hydrogen from the purge stream, and ammonia can be recovered by absorption with water.

AMMONIA SYNTHESIS CATALYSTS Promoted iron catalysts are typically used for ammonia synthesis. The catalysts are supplied as magnetite (Fe3O4) and reduced in the reactor to produce metallic iron: Fe3 O4 ðsÞ þ 4H2 ðgÞ/3FeðsÞ þ 4H2 OðgÞ

R-11.3

Removal of oxygen from the solid magnetite pellet produces a porous iron pellet, with about 10 to 20 m2/g surface area. The catalyst pellets are promoted with small quantities of Al2O3 and K2O.

Ammonia and Derivatives

Oxygen compounds, such as H2O, CO, and CO2, behave as temporary poisons. For example, water in the feed gas would reverse R-11.3, producing non-catalytic iron oxide instead of catalytic iron metal. Once the poison is removed, the iron oxide is gradually reduced, and the catalytic activity is restored. Permanent poisons include sulfur compounds, which can be introduced in the feed gas or by compressor oil mists. Consequently, ammonia plants tend to use oil-free compressors. Ruthenium catalysts are much more active than the traditional iron catalyst.3 This allows the reactor to run at lower temperatures and pressures, which result in a considerable savings in compression and pressure vessel costs. On the other hand, ruthenium is a very rare and expensive metal, and it is not clear that the compression and pressure vessel cost savings justify the high catalyst cost.

AMMONIA SYNTHESIS KINETICS The overall rate of ammonia synthesis with an iron catalyst can be written4 as: kr ¼ khIP hTP

Eqn. 11.2

Where kr ¼ overall rate of ammonia synthesis, k ¼ uninhibited surface reaction rate, hIP ¼ effectiveness factor due to intra-particle pore diffusion restrictions, hTP ¼ effectiveness factor due to temporary poisons. The surface reaction rate is most commonly modeled using the Temkin-Pyzhev rate equation. This equation is based on the following assumptions: (1) N2 adsorption is rate-limiting; (2) the N2 adsorption rate is affected by pre-adsorbed molecules on the surface; (3) the reaction is reversible. The rate equation is based on the gas fugacities, rather than the partial pressures, to account for the non-ideal gas behavior at high pressure. The fugacity of a perfect gas is its partial pressure. A fugacity coefficient is used to account for the non-ideal behavior of the gas. fi ¼ vi Pi

Eqn. 11.3

Where vi ¼ fugacity coefficient of gas species i, Pi ¼ partial pressure of gas species i. The rate expression is: 2 3 a 2 ð1aÞ vNH3 fNH3 fH EA 1 1 k ¼ ko exp KNH3 fNH3 2 2 3 R T To vN2 vH2 fNH3 fH32 Eqn. 11.4

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Ammonia and Derivatives

The effectiveness factor due to temporary catalyst poisons, hTP, can be written4 as: pH2 2 0:5 Eqn. 11.5 hTP ¼ 1 þ CPw pNH3 where Pw is the partial pressure of oxygen compounds expressed as an equivalent water partial pressure: Pw ¼ PH2 O þ PCO þ 2PCO2 þ /

Eqn. 11.6

The rate constant in Eqn. 11.5 is given as: 6

C ¼ ð1:78X10

8; 700 atmÞ exp T

Eqn. 11.7

where T is the temperature in Kelvin. The rate-limiting mass transfer step during ammonia synthesis is diffusion of ammonia out of the catalyst. For a spherical catalyst, the effectiveness factor due to intraparticle pore diffusion, hIP, is given by: 3 1 1 Eqn. 11.8 hIP ¼ f tanh ðfÞ f A plot of hIP versusf, the Thiele modulus, can be found in Figure 3.8. The relationship between hIP and f for non-spherical catalysts can be found in heterogeneous catalysis textbooks. For ammonia synthesis, the Thiele modulus is defined as: sﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃﬃ jDS k Eqn. 11.9 f ¼ 2 DNH3 The equivalent catalyst particle diameter, Ds, is given by: Ds ¼

6X ðcatalyst particle volumeÞ ðcatalyst particle external surface areaÞ

Eqn. 11.10

For a spherical catalyst particle, Ds is the true particle diameter. The shape factor, j, is one for spherical particles, and less than one for non-spherical particles. The pseudo-first-order reaction rate, k), is defined as the derivative of the reaction rate uninhibited by mass transfer with respect to the ammonia mole fraction: k) ¼

d ðkhTP Þ dyNH3

Note that k is given by Eqn. 11.4 and hTP is given by Eqn. 11.5.

Eqn. 11.11

Ammonia and Derivatives

The effective diffusion coefficient for ammonia, DNH3 , is given by: DNH3

1:5 3 T 1 ¼ s 273 p

1 yNH3 s yH2 s yN 2 s þ DNH3 H2 DNH3 N2

Eqn. 11.12

Where 3 ¼ catalyst particle void volume, typically 0.55, s ¼ catalyst tortuosity factor, typically 2.2, T ¼ temperature, Kelvin, P ¼ pressure, atmospheres, yNH3 s ¼ ammonia mole fraction at the catalyst surface, yH2 s ¼ hydrogen mole fraction at the catalyst surface, yN2 s ¼ nitrogen mole fraction at the catalyst surface, DNH3 H2 ¼ binary ammonia-hydrogen diffusion constant, typically 0.257 m2/h, DNH3 H2 ¼ binary ammonia-nitrogen diffusion constant, typically 0.0799 m2/h. Gramatica and Pernicone4 show that the effectiveness factor due to intra-particle pore diffusion restrictions, hIP , is about 0.2 near the reactor inlet, where the driving force for reaction is the greatest; and about 0.9 near the reactor outlet, where the driving force for force for reaction is the smallest.

AMMONIA AS NITROGEN FERTILIZER A substantial fraction of produced ammonia is applied directly to farm fields as anhydrous ammonia. Farmers drag an injection rake behind a tractor, and inject liquid ammonia a few cm below the soil surface. As the liquid ammonia vaporizes it permeates the soil. Most of the remaining ammonia is converted into other types of nitrogen fertilizers, including the solids urea, ammonium nitrate, ammonium phosphates, and ammonium sulfate, as well as liquid ureaeammonium nitrate (UAN) solutions. Ammonium nitrate is produced by neutralizing nitric acid with ammonia. Both the ammonium and nitrate ions provide nitrogen to the plants. The production of urea and nitric acid from ammonia will be described in greater detail later in this chapter. The ammonium phosphates, mono-ammonium phosphate (MAP) and di-ammonium phosphate (DAP), are the most common phosphate fertilizers. These are made by neutralizing an impure wet-process phosphoric acid with ammonia. These fertilizers contain two primary nutrients, nitrogen and phosphorus, as well as numerous secondary and micro nutrients due to the impurities in the phosphoric acid. Ammonium sulfate can be made by neutralizing sulfuric acid with ammonia, but much of it is a byproduct from caprolactum production, which is the monomer used to produce Nylon-6. The Powerspan process5,6 for SO2 (or SO2 plus CO2) removal from flue gas also produces ammonium sulfate (see Chapter 6). Ammonium sulfate provides

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Ammonia and Derivatives

a primary nutrient, nitrogen, and a secondary nutrient, sulfur. As a salt of a strong acid, sulfuric acid, and a relatively weak base, ammonia, ammonium sulfate forms acidic solutions. Consequently, it is not recommended for acid soils, but it is widely used to neutralize alkaline soils. UAN solutions are made by dissolving urea and ammonium nitrate in water. These are foliar fertilizers that can be applied to growing plants through the irrigation system.

AMMONIA AS A TRANSPORTATION FUEL Ammonia could be used as a transportation fuel, either by burning it in an internal combustion engine, or by feeding it to a fuel cell. In either case, the reaction products are nitrogen and water: 2NH3 ðgÞ þ 1:5O2 ðgÞ/N2 ðgÞ þ 3H2 OðgÞ

DH rxn ¼ 633:6 kJ=gmole R-11.4

Like hydrogen, the combustion of ammonia produces no greenhouse gases. Unlike hydrogen, ammonia, with a 33 oC boiling point, is easily liquefied. The heat of combustion of ammonia is 18.6 MJ/kg, versus about 44 MJ/kg for gasoline. The US Department of Energy target for hydrogen storage is 5 wt.% hydrogen, which corresponds to only 6 MJ/kg. Compared to hydrogen, ammonia is a much more energydense fuel. There are disadvantages. First of all, converting hydrogen to ammonia and then burning the ammonia produces only 87% of the heat that would be obtained if the hydrogen were burned. The difference is due to the heat of formation of ammonia, shown in R-11.1. Like most fuels, ammonia is flammable. Ammonia is also toxic at high concentrations, primarily due to its alkaline nature. As a compressed fuel, a small crack in a fuel line could produce a large leak. Ammonia was once commonly used as a refrigerant in household appliances, but it was replaced by safer chlorofluorocarbon refrigerants. Aqueous ammonia is sold as a 30 wt.% solution. This would be safer to handle, since it could be stored at atmospheric pressure, but it would have a lower energy density.

UREA The structure of urea is shown in Figure 11.3. Urea hydrolyzes in soil to form ammonia and carbon dioxide. Compared to other fertilizers, the release of ammonia is more gradual, so there is a reduced tendency to “burn” the plant. O H H N-C-N H H

Figure 11.3 Structure of urea.

Ammonia and Derivatives

The feedstocks for urea production are ammonia and carbon dioxide. When syngas is converted to hydrogen for ammonia production, carbon dioxide is a byproduct (see Figure 7.2), and this is the usual source of carbon dioxide for urea production. In a typical urea production plant,7 ammonia and carbon dioxide, typically at a 3e4 molar ratio, are fed to a reactor operating at 14e18 MPa and about 190 oC. There, ammonia and carbon dioxide initially react to form ammonium carbamate. 2NH3 ðgÞ þ CO2 ðgÞ4NH2 Cð¼ OÞ2 NH4 ðaqÞ

R-11.5

A portion of the ammonium carbamate decomposes in the reactor to form urea and water: NH2 Cð¼ OÞ2 NH4 ðaqÞ/NH2 C ¼ ONH2 ðaqÞ þ H2 Oð[Þ

R-11.6

The reactor effluent solution is then sent to a stripping section where the solution is heated, sometimes in the presence of a stripping gas, to remove unreacted ammonia and further decompose the ammonium carbamate via R-11.6. The urea/water solution is concentrated in an evaporator to about 75% urea. This urea solution, along with air, is fed to a fluidized bed granulator to produce solid urea granules. In addition to fertilizer use, urea is used as an animal feed supplement. Increasing quantities of urea are being used to reduce NOx emissions from diesel engines. As shown in Chapter 2, NOx emissions from power plants are reduced by feeding the flue gas and ammonia to a selective catalytic reactor (SCR), where NOX compounds and ammonia react to form N2 and water. Urea is used in place of ammonia for this purpose in automobiles and trucks, as urea is more conveniently stored in a vehicle.

NITRIC ACID AND AMMONIUM NITRATE A mixture of ammonia and air is burned8 over a 90% platinum, 10% rhodium gauze catalyst to produce nitric oxide: 4NH3 ðgÞ þ 5O2 ðgÞ/4NOðgÞ þ 6H2 OðgÞDHrxn ¼ 906:0 kJ$gmole

R-11.7

This reaction is deliberately run in a mass-transfer limited regime. If it were operated in a more kinetically limited regime, the reaction products would be primarily N2 and N2O. In a typical process, 9 the ammonia and air feed mixture is at 0.7 to 1 MPa, and is preheated to 170e180 oC. The ammonia concentration in the feed mixture is about 10%. The gauze temperature reaches 920e940 oC. Hot gases leaving the burner are cooled and fed to an absorber/oxidizer.10 Nitric acid reacts homogeneously with oxygen to yield nitrogen dioxide: 2NOðgÞ þ O2 ðgÞ/2N2 OðgÞ

DHrxn ¼ 114:4 KJ$gmole

R-11.8

349

350

Ammonia and Derivatives

The nitrogen dioxide then dimerizes to form dinitrogen tetroxide: 2N2 OðgÞ/N2 O4 ðgÞ

DHrxn ¼ 57:12 KJ$gmole

R-11.9

Dinitrogen tetroxide is absorbed by water to form nitric and nitrous acids: 2N2 O4 ðgÞ þ H2 OðlÞ/2HNO3 ðaqÞ þ 2HNO2 ðaqÞ

R-11.10

Nitrous acid decomposes to yield nitrogen dioxide and nitric acid: 2HNO2 ðaqÞ/NO2 ðgÞ þ NOðgÞ þ H2 OðlÞ

R-11.11

The nitrogen dioxide and nitric oxide liberated by (R-11.11) react to form additional nitric acid via (R-11.8), (R-11.9), and (R-11.10). Ammonium nitrate is made by neutralization of nitric acid with ammonia. There are two primary uses of ammonium nitrate. The first is as a nitrogen fertilizer. Both the ammonium and the nitrate ions provide nitrogen to plants. The second use is as an explosive. The nitrate ion is an oxidizer, and the ammonium ion is a fuel. In pure ammonium nitrate, the oxidizing potential of the nitrate ion is greater than the reducing potential of the ammonium ion; so explosives are formulated by mixing ammonium nitrate with a fuel oil (ANFO). This is a popular explosive in mining and construction, primarily due to its low cost. When coal is strip-mined, the first deposits mined are generally those with the lowest stripping ratios (overburden-to-coal ratio). As those deposits are depleted, mining moves toward deposits with higher stripping ratios, meaning that more overburden must be removed for every ton of coal mined. The general procedure is to blast the rock, and then remove it. Consequently, the demand for ANFO is increasing at a faster fractional rate than the demand for coal. There are two common grades of ammonium nitrate. The low density grade is intended for ANFO.11 The higher surface area of the low density ammonium nitrate allows for more intimate contact between the ammonium nitrate and the fuel oil, which leads to better explosive performance. Fertilizer grade ammonium nitrate is a high density product. Although less-than-optimum, fertilizer grade ammonium nitrate can also be used in ANFO. The agricultural use of ammonium nitrate faces increased regulation due to the potential use of ANFO as explosives. Ammonium nitrate is also used as a component in UAN (ureaeammonium nitrate) solutions. This is an aqueous solution that may be applied to plants through the irrigation system. Solid fertilizers must normally be applied before crops are planted. UAN, on the other hand, may be applied throughout the growing season.

FUTURE OF NITROGEN FERTILIZERS Agricultural productivity greatly increased in the last 100 years, due largely to the use of synthetic fertilizers, pesticides, and hybrid crop varieties. Increased productivity led to

Ammonia and Derivatives

a decline in the prevalence of famine. The larger food supply allowed the world’s population to grow. In 1900, a few years before Fritz Haber discovered ammonia synthesis, the world’s population was about 1.6 billion people. By 2008, the population more than quadrupled to about 6.7 billion people. There is concern that continued increases in agricultural productivity may not be able to keep up with our growing population, and famine may become common again. In the developed world, organic foods are popular. These are foods grown without the aid of synthetic fertilizers or pesticides. A relatively small but substantial portion of our food supply now comes from organic farms. While successful on a limited scale, organic farming on a wide scale may not be able to produce sufficient food to feed our large and growing population. Organic foods are substantially more expensive than foods grown with synthetic fertilizers and pesticides. This is a minor inconvenience to affluent consumers in developed nations, but it makes organic foods unaffordable to poor consumers. We may need to use synthetic fertilizers and pesticides for a very long time. A number of commenters decry the dependence of modern farming on fossil fuels, particularly petroleum. Nitrogen fertilizers are a large component of the petroleum consumption of a typical farm. When the word “petroleum” is used, however, most readers assume that this means crude oil, or crude oil derived fuels such as diesel used by farm machinery. The commenter has usually failed to mention that nitrogen fertilizers are usually made from natural gas, which is far more abundant than crude oil. If natural gas prices rise, nitrogen fertilizers can be made from coal, or alternatively, hydrogen produced by nuclear power or renewable energy. A crude oil shortage may lead to a shortage of fuel for farm machinery, but there will be little long-term impact on nitrogen fertilizer production.

REFERENCES 1. Jennings JR, Ward SA. Ammonia synthesis. In: Twigg MV, ed. Catalyst Handbook. 2nd ed. Wolfe Publishing; 1989. p. 384-440. 2. Riddler DE, Twigg MV. Steam reforming. In: Twigg MV, ed. Catalyst Handbook. 2nd ed. Wolfe Publishing; 1989. p. 225-2285. 3. Kowalczyk Z, Jodzis S, Sentek J. Studies on kinetics of ammonia synthesis over ruthenium catalyst supported on active carbon. Appl Cattal A: Gen. 1996;138:83-91. 4. Gramatica G, Pericone N. Kinetics of ammonia sythesis and influence on converter design. In: Jennings JR, ed. Catalytic Ammonia Synthesis: Fundamentals and Pracitice. Plenum Press; 1991. 5. Powerspan, <www.powerspan.com>. 6. Pennline H, Richards G, Plasynski SI. Ammonia-based process for multicomponent removal from flue gas. <www.netl.doe.gov/publications/factsheets/rd/R&;D043.pdf>. 7. Mavrovic I, Shirley Jr AR, Coleman GR. Urea. In: Kirk-Othmer Encyclopedia of Chemical Technology. 4th ed. John Wiley & Sons; 1998. Supplemental Volume. 8. Andrew SPS. The oxidation of ammonia. In: Keleti C, ed. Nitric Acid and Fertilizer Nitrates. Marcel Dekker; 1985. 9. Honti GD. Various processes for the production of commercial-grade nitric acid. In: Keleti C, ed. Nitric Acid and Fertilizer Nitrates. Marcel Dekker; 1985.

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10. Andrew SPS. The absorption and oxidation system. In: Keleti C, ed. Nitric Acid and Fertilizer Nitrates. Marcel Dekker; 1985. 11. Dyno Nobel, Inc. Ammonium nitrate, LoDANÔ industrial grade. <www.dynonobel.com/NR/ rdonlyres/C6A0B2EA-9B00-40EC-9096-B5ABE94A67D8/0/AmmoniumNitrate_LoDan10507. pdf>.

CHAPTER

12

Methanol and Derivatives Contents Reaction Chemistry and Catalysts Methanol Synthesis Equilibria Methanol Synthesis Flowsheet Methanol Synthesis Kinetics Methanol as a Transportation Fuel Dimethyl Ether Two Pot Synthesis of DME One Pot Synthesis of DME DME to Hydrocarbons: ExxonMobil MTG Process Methanol to Hydrocarbons: UOP/HYDRO MTO Process Methanol and DME to Hydrocarbons: Future Prospects References

353 354 356 357 359 359 360 361 365 368 370 370

REACTION CHEMISTRY AND CATALYSTS At high pressure, in the presence of a catalyst, hydrogen, and carbon monoxide may react to form methanol: o ¼ 94:5 kJ=gmole 2H2 ðgÞ þ COðgÞ4CH3 OH DHrxn

R-12.1

The syngas stream will usually contain a small quantity of carbon dioxide, which may also react with hydrogen to form methanol plus water. o 3H2 ðgÞ þ CO2 ðgÞ4CH3 OHðgÞ þ H2 OðgÞ DHrxn ¼ 53 kJ=gmole

R-12.2

Methanol synthesis catalysts also have water gas shift activity, which may account for their ability to synthesize methanol from either carbon monoxide or carbon dioxide. o ¼ 41:21 kJ=gmole COðgÞ þ H2 OðgÞ4CO2 ðgÞ þ H2 ðgÞ DHrxn

R-12.3

A potential problem with methanol synthesis is that there are a number of potential alternative and side reactions that may also occur. For example, crude methanol often contains small quantities, typically1 100 ppm, of dimethyl ether produced by the following reaction: o ¼ 16 kJ=gmole 2CH3 OHðgÞ4CH3 OCH3 ðgÞ þ H2 OðgÞ DHrxn Coal Gasification and Its Applications. ISBN B978-0-8155-2049-8.10012-9, doi:10.1016/B978-0-8155-2049-8.10012-9

R-12.4

Ó 2011 Elsevier Inc. All rights reserved.

353

354

Methanol and Derivatives

Carbon monoxide and hydrogen can also react to form higher alcohols (shown here as ethanol), methane, or higher hydrocarbons via Fischer-Tropsch synthesis. o ¼ 255 kJ=gmole 4H2 ðgÞ þ 2COðgÞ/CH3 CH2 OHðgÞ þ H2 OðgÞ DHrxn R-12.5 o ¼ 216:2 kJ=gmole 3H2 ðgÞ þ COðgÞ/CH4 ðgÞ þ H2 OðgÞ DHrxn

R-12.6

o varies ð2n þ 2ÞH2 ðgÞ þ nCOðgÞ/HðCH2 Þn Hð1 þ gÞ DHrxn

R-12.7

Thus a key requirement for methanol catalysts is that they must be selective. Methanol was first commercially produced from syngas in 1923. Before the 1960s, the most commonly used catalyst was a co-precipitated mixture of ZnO and Cr2O3. In this catalyst, zinc is the active species, and chromia prevents the zinc from sintering. The reactor typically operated at 35 MPa and 350 to 400 C. Methanol synthesis catalyst development up to 1955 was reviewed by Natta.2 In 1963 Imperial Chemical Industries (ICI) introduced a much more active catalyst, a co-precipitated mixture of CuO, ZnO, and Al2O3. Copper is the active catalytic species, and alumina inhibits catalyst sintering. This catalyst is much more active, which allowed a reduction in reactor pressure to about 5 to 10 MPa. This will result in a considerable cost savings. The higher activity of the catalyst also allowed a reduction of typical reactor temperatures to about 200 to 250 C . In addition to improving the reaction equilibrium, this lower temperature was necessary to prevent catalyst sintering. Essentially all modern methanol synthesis catalysts are variations of the ICI catalyst.

METHANOL SYNTHESIS EQUILIBRIA An equilibrium constant for the synthesis of methanol from carbon monoxide (R-12.1) can be written as: KMeOHCO ¼

PCH3 OH PCO ðPH2 Þ2

Eqn. 12.1

Calculated values for this equilbrium constant as a function of temperature and pressure are shown in Figure 12.1. Since three moles of gas react to produce one mole of gas, higher pressures push the reaction to the right. Higher reaction temperatures reduce the equilibrium methanol concentration. Methanol can also be made from carbon dioxide via R-12.2. For this reaction, an equilibrium constant may be written as: KMeOHCO2 ¼

PCH3 OH PH2 O PCO ðPH2 Þ3

Eqn. 12.2

Methanol and Derivatives

Equilibrium Constant

100

7 MPa

10

10 MPa

1 5 MPa

0.1

0.01 200

210

220

230

240

250

260

270

280

290

300

Temperature, C

Figure 12.1 Methanol synthesis equilibria based on a carbon monoxide feedstock (R-12.1) and using the equilibrium constant shown in Eqn. 12.1. The units of the equilibrium constant are MPa2. A stoichiometric 2:1 hydrogen to carbon monoxide feed was assumed.

Calculated values for this equilibrium constant are shown in Figure 12.2. Like methanol synthesis from carbon monoxide, the number of moles of gas are reduced as the reaction proceeds, so high pressures will push the reaction to the right. Again, the equilibrium constant declines with increasing temperature. The equilibrium calculations show that the reverse water gas shift reaction (R-12.3) is a competing reaction, and becomes especially prominent as the temperature increases. 1

Equilbrium Constant

10 MPa 7MPa

0.1

5 MPa

0.01

0.001 200

210

220

230

240

250

260

270

280

290

300

Temperature, C

Figure 12.2 Methanol synthesis equilibria based on a carbon dioxide feedstock (R-12.2) and using the equilibrium constant shown in Eqn. 12.2. The units of the equilibrium constant are MPa2. A stoichiometric 3:1 hydrogen to carbon monoxide feed ratio was assumed. The reverse water gas shift reaction (R-12.3) is also signiﬁcant at these conditions.

355

Methanol and Derivatives

Percent Conversion of CO or CO2 at 7 MPa

356

100 90 Methanolfrom CO

80 70 60 50

Methanol from CO2

40 30 20

CO from CO2

10 0 200

210

220

230

240

250

260

270

280

290

300

Temperature, C

Figure 12.3 Equilbrium synthesis of methanol at 7 MPa from either a carbon monoxide feedstock (R-12.1) or a carbon dioxide feedstock (R-12.2). When a carbon dioxide feedstock is used, the reverse water gas shift reaction (R-12.3) is signiﬁcant, so a portion of the carbon dioxide is converted to carbon monoxide. For the carbon monoxide feedstock, a stoichiometric 2:1 hydrogen to carbon monoxide feed ratio was used. For the carbon dioxide feedstock, a stoichiometric 3:1 hydrogen to carbon dioxide feed ratio was used.

The ability to synthesize methanol from carbon dioxide raises the intriguing possibility that carbon dioxide, normally considered a waste product, can be converted to a useful product. This would be practical only if a large volume of hydrogen were available from a non-fossil fuel source. For example, hydrogen could be made by the electrolysis of water using off-peak electric power from a nuclear power plant. Figure 12.3 compares the equilibrium synthesis of methanol from either carbon monoxide or carbon dioxide. The driving force for reaction R-12.2, synthesis from carbon dioxide, is smaller than the driving force for synthesis from carbon monoxide, R-12.1, leading to a lower per-reactor pass conversion. Also, when carbon dioxide is fed to the reactor, the reverse water gas shift reaction, R-12.3, becomes significant, especially at higher reactor temperatures.

METHANOL SYNTHESIS FLOWSHEET On a qualitative level, the synthesis of methanol from syngas is much like the synthesis of ammonia from a hydrogen/nitrogen mixture. Conversion of the feed gas to product is substantially less than complete, so the product is separated from the reactor effluent by condensation. The unconverted gas is recompressed and recycled. A portion of the recycled gas is purged to prevent the accumulation of inert gases. High pressures give favorable equilibria, but high pressure equipment is expensive, so the operating pressure is a compromise. Likewise, low temperatures give good equilibria, but higher

Methanol and Derivatives

synthesis reactor

feed gas 200 °C

quench gas purge gas cooling water 250 °C

crude methanol

recycle gas

fresh feed gas

feed compresor

recycle compressor

Figure 12.4 Methanol synthesis ﬂowsheet based on a quench reactor.

temperatures give faster reaction rates, so the operating temperature range is also a compromise. The gas fed to the reactor is heated to attain the desired reaction rate, but the reactor is cooled to remove the exothermic heat of reaction. A flow sheet based on a quench reactor design is shown in Figure 12.4. Note the similarity of this flowsheet to the ammonia synthesis flow sheet shown in Figure 11.2. Unlike ammonia, the produced methanol contains substantial levels of impurities. If the feed gas contains carbon dioxide, the crude methanol will contain water. The crude methanol will also typically1 contain about 100 ppm dimethyl ether, 20 ppm carbonyl compounds, and 400 ppm higher alcohols (C2þ) due to less-than-perfect catalyst selectivity. Distillation is normally used to purify the crude methanol product to meet specifications.

METHANOL SYNTHESIS KINETICS Methanol synthesis kinetics have been investigated by a number of research teams. The kinetic expressions reported here are from Graaf et al.,3 who gathered laboratory reaction data on a commercial CuO/ZnO/Al2O3 catalyst, Haldor Topsoe Mk 101, over a 215e245 C temperature range and a 15e50 bar (1.5e5 MPa) pressure range. They considered a number of alternative kinetic expressions, and then selected those that best

357

358

Methanol and Derivatives

fit their data. Three reactions were significant, the synthesis of methanol from carbon monoxide (R-12.1), the synthesis of methanol from carbon dioxide (R-12.2), and the water gas shift reaction (R-12.3). The kinetic expression for the synthesis of methanol from carbon monoxide is: f 3=2 k1 KCO fCO fH2 2=2 CH3 OH fH2 KMeOHCO rMeOHCO ¼ Eqn. 12.3 1=2 1 þ KCO fCO þ KCO2 fCO2 fH2 þ k2 fH 2O Where rMeOH-CO ¼ rate of methanol synthesis from carbon monoxide per weight of catalyst, mole/(s-kg), Þ, k1 ¼ ð2:69 0:14Þ 107 expð109;900200 RT R ¼ gas constant, 8.314 J/mole-K, T ¼ temperature, K, KCO ¼ ð7:99 1:28Þ 107 expð58;100600 Þ, RT fi ¼ fugacity of species i, bar, KMeOH-CO ¼ equilibrium constant for methanol synthesis from carbon monoxide, bar2 (note: the values shown in Figure 12.1 are in MPa2, so they need to be divided by 100 to convert them to bar2). 67; 400 600 7 KCO2 ¼ ð1:02 0:16Þ 10 exp RT k2 ¼ ð4:13 1:51Þ 10

11

104; 500 1; 100 exp RT

The kinetic expression for the synthesis of methanol from carbon dioxide (R-12.2) is: fCH3 OH fH2 O 3=2 f k3 KCO2 CO2 fH2 3=2 fH2 KMeOHCO2 Eqn. 12.4 rMeOHCO2 ¼ 1=2 1 þ KCO fCO þ KCO2 fCO2 fH2 þ k2 fH2 O Where rMeOH-CO2 ¼ rate of methanol synthesis from carbon dioxide per weight of catalyst, mole/(s-kg). 65; 200 200 2 k3 ¼ ð4:36 0:25Þ 10 exp RT KMeOH-CO2 ¼ equilibrium constant for methanol synthesis from carbon dioxide, bar2 (note:

the values shown in Figure 12.2 are in MPa2, so they need to be divided by 100 to convert them to bar2).

Methanol and Derivatives

The kinetic expression for the water gas shift reaction is: fH2 O fCO k4 KCO2 fCO2 fH2 KWG rWG ¼ 1=2 1 þ KCO fCO þ KCO2 fCO2 fH2 þ k2 fH2 O

Eqn. 12.5

Where rWG ¼ rate of the water gas shift reaction per weight of catalyst, mole/(s-kg). 123; 400 1; 600 8 k4 ¼ ð7:31 4:90Þ 10 exp RT KWG ¼ equilibrium constant for the water gas shift reaction (see Chapter 7 for values). In a further investigation, Graaf, et al.4 showed that mass transfer restrictions have a significant impact on the rate of reaction at temperatures greater than about 245 C. At 10 bar (1 MPa), a relatively low synthesis pressure, the gas mean free path is long compared to typical catalyst pore diameters, and Knudsen diffusion predominates. At 100 bar (10 MPa), the gas mean free path is shorter and bulk diffusion predominates.

METHANOL AS A TRANSPORTATION FUEL Methanol is a liquid at ambient conditions (boiling point is 64.7 C), and has been widely considered as a transportation fuel. The Chinese government recently adopted an M85 fuel standard, which nominally consists of 85% methanol and 15% hydrocarbons. The hydrocarbon content is adjusted so that the blended fuel meets gasoline volatility specifications. Methanol, especially M85, can be burned in engines designed to burn gasoline, but methanol may be corrosive to materials designed to contact gasoline. Methanol is also hydrophilic, so it may absorb water, which would have a negative impact on its performance. The primary drawback to methanol as a fuel is that it is toxic. The OSHA permissible exposure limit is 200 ppm. Leaked methanol can contaminate water supplies, but it is rapidly biodegraded. The heat of combustion for methanol, at 22.3 MJ/kg, is substantially less than gasoline, typically 44 MJ/kg, so a vehicle burning methanol will have a significantly lower range compared to gasoline.

DIMETHYL ETHER Dimethyl ether (DME) is formed by the dimerization of methanol: 2CH3 OHðgÞ4CH3 OCH3 ðgÞ þ H2 OðgÞ DH orxn ¼ 16 kJ=gmole

(R-12.4)

DME is a gas at ambient conditions, with a 25 C boiling point and a 0.5 MPa vapor pressure at 20 C. DME is slightly polar, and is nearly nontoxic.5

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Methanol and Derivatives

DME has a very promising future, either as a fuel, or as an intermediate in the production of hydrocarbon fuels and chemicals. Currently, a primary use of DME is as a propellant in personal care products, such as hairspray. One of the uses for DME is as a substitute for liquefied petroleum gas (LPG), which is primarily a blend of propane and butane. The boiling point of DME is between that of propane, 42 C, and butane, 0.5 C for normal butane, 12 C for isobutane, so liquid DME can be stored in tanks designed for LPG, so long as the proper polymers are chosen for seals. LPG is a popular fuel for home heating and cooking, especially in regions that lack household distribution of natural gas. In the USA, LPG, primarily propane, is used in rural areas. In many other countries, bottled LPG is also a popular cooking fuel in urban areas. Since LPG is produced as a by-product of natural gas production or oil refining, production is often not sufficient to meet demand. There are a number of DME production plants either operating or under construction in China.6 DME from these plants is sold primarily as a cooking fuel. Since DME contains an oxygen atom, the gas phase heat of combustion, 1,460 kJ/mole, is less than propane, 2,219 kJ/mole, so a greater quantity must be burned to generate the same amount of heat. DME is also being considered as a substitute fuel for diesel engines. The cetane index for DME is 55, and, compared to conventional diesel fuel, DME produces about 90% less NOx emissions and essentially no particulates.7 Diesel automobiles are popular in Europe because they are more fuel efficient than gasoline automobiles. In the USA, diesel automobiles are uncommon due to the difficulty of meeting US specifications on NOx and particulate emissions. As a fuel for diesel engines, DME may enable wider adoption of diesel engines in US automobiles. DME is widely touted as a means of making a diesel fuel substitute from biomass. Biodiesel is made from plant or animal fats and oils. This places biodiesel in direct competition with food production. DME, on the other hand, can be made from syngas produced by gasifying cellulosic biomass or coal, so DME is not a direct competitor to food production. Syngas can also be converted to diesel via FischerTropsch synthesis, as described in the next chapter, but DME synthesis is considerably less complex. The energy density of DME is 31.7 MJ/kg, compared to about 45 MJ/kg for diesel. The primary challenges in converting diesel vehicles to DME is changing the fuel tank and the fuel distribution system.

TWO POT SYNTHESIS OF DME There are two approaches to DME synthesis. The most common current approach is the two pot synthesis, in which methanol is made from syngas, and then the methanol is

Methanol and Derivatives 30

Equlibrium Constant

25 20 15 10 5 0 200

250

300

350

400

Temperature, C

Figure 12.5 Equilibrium conversion of methanol into dimethyl ether, R-12.4, using the equilibrium constant shown in Eqn. 12.6, at 1 MPa.

converted to DME via R-12.4. An equilibrium constant for DME production can be written as: KDME ¼

PDME PH2 O ðPCH3 OH Þ2

Eqn. 12.6

This equilibrium constant, at 1 MPa, is shown in Figure 12.5 as a function of temperature. Since R-12.4 converts two moles of gas reactants into the same number of moles of products, the reaction pressure has only a small effect on the equilibrium. Figure 12.6 shows the Topsoe DME synthesis process.8 Methanol is preheated and fed to a reactor at 260 to 300 C . The reactor is packed with a modified alumina catalyst. The reactor is operated adiabatically, and the temperature of the gases passing through the reactor rises to about 375 to 400 C, due to the exothermic heat of reaction. The reactor effluent is cooled and sent to the DME column. The overhead condensate from this column is the DME product. Bottoms from the DME column is fed to the water column, and by-product water is produced as the bottoms from the water column. The overhead product from the water column contains unconverted methanol, and this stream is fed to the offgas absorber to scrub uncondensed gasses from the DME column. The bottoms stream from the absorber is recycled to the feed stream.

ONE POT SYNTHESIS OF DME In the one pot synthesis of DME, a newer approach, the reactor contains both the methanol synthesis catalyst and a solid acid catalyst for the conversion of methanol to DME. Two approaches to catalyst formulation are taken. One is to load methanol

361

362

Methanol and Derivatives

off gas methanol

off gas absorber recycle

DME product

DME column water column reactor

waste water

Figure 12.6 Topsoe DME synthesis process (two pot synthesis).8

synthesis catalyst onto a solid acid support. The second is to simply mix methanol synthesis and solid acid catalysts. The one pot approach offers two major advantages over the two pot approach. First of all, the process has fewer unit operations, which implies a lower capital cost. The second is that the conversion of methanol to DME pushes the methanol synthesis reaction to the right, meaning that there is a greater overall per pass conversion of syngas to products. One might think, based on R-12.1 and R-12.4, that syngas might react to form a DME/water mixture. At typical reactor conditions, however, the right side of the water gas shift reaction (R-12.3) is favored; so most of the water produced by R-12.4 reacts with carbon monoxide in the feed gas to produce carbon dioxide and hydrogen. A more realistic description of the one pot synthesis of DME is to combine R-12.1, R12.3, and R-12.4 to give: o 3COðgÞ þ 3H2 ðgÞ4CH3 OCH3 ðgÞ þ CO2 ðgÞ DHrxn ¼ 243 kJ=gmole R-12.5

Reactions R-12.1, R-12.3, and R-12.4 are all exothermic, so the combination of these three reactions, R-12.5, is strongly exothermic. This pushes the reaction to the right. Figure 12.7 shows that the equilibrium per-pass conversion of carbon monoxide to

Methanol and Derivatives

Figure 12.7 Single pass equilibrium conversion of carbon monoxide at 7 MPa. Methanol synthesis is according to reaction R-12.1 using a stoichiometric 2:1 hydrogen to carbon monoxide feed. Dimethyl ether synthesis was by reaction R-12.5 using a stoichiometric 1:1 hydrogen to carbon monoxide feed.

DME is greater than the per-pass conversion to methanol at the same reaction conditions. Water and methanol are produced as by-products, but at low concentrations. In the calculations used to produce Figure 12.7, less than 1.31% of the carbon monoxide was converted to methanol, and less than 0.63% of the hydrogen was converted to water. Just as it is possible to synthesize methanol from carbon dioxide and hydrogen, it is also possible to synthesize DME from carbon dioxide and hydrogen. If the only source of carbon fed to the reactor is carbon dioxide, then the water gas shift reaction cannot move to the right, and a combination of R-12.2 and R-12.4 gives: o 2CO2 ðgÞ þ 6H2 ðgÞ4CH3 OCH3 ðgÞ þ 3H2 OðgÞ DHrxn ¼ 119:5 kJ=gmole R-12.6

The driving force for this reaction is much smaller, so, as shown in Figure 12.8, the equilibrium per pass conversion of carbon dioxide is much less. Methanol is a significant by-product, which could be recycled to give further DME. The reverse water gas shift reaction, R-12.2, becomes significant at high reaction temperatures. The highly exothermic nature of R-12.5 can lead to hot spots within the reactor. With increasing temperature, DME may further dehydrate on the solid acid catalyst to form hydrocarbons: o varies nCH3 OCH3 ðgÞ/hydrocarbonsðgÞ þ nH2 OðgÞ DHrxn

R-12.7

These hydrocarbons may be the desired product, but hydrocarbon formation is also usually accompanied by coke formation on the catalyst. In other catalytic systems, coke is removed from catalyst by burning it off under controlled conditions. For the one pot

363

Methanol and Derivatives

60

Equilbrium Conversion of Carbon Dioxide,

364

50 carbon dioxide

40 dimethyl ether

30

20 methanol

10 carbon monoxide

0 210

220

230

240

250

260

270

280

290

300

Temperature, C

Figure 12.8 Equilbrium conversion of carbon dioxide at 7 MPa to dimethyl ether, methanol, and carbon monoxide. All values are based on the percentage of carbon dioxide converted, assuming a stoichiometric 3:1 hydrogen to carbon dioxide feed ratio based on R-12.6.

synthesis of DME, burning off the coke in this fashion would destroy the methanol synthesis catalyst. Consequently, much of the research on the one pot synthesis technique has focused on catalysts and reaction systems that give high DME synthesis activity but avoid hydrocarbon formation. Most of the investigated catalyst systems use a conventional CuO/ZnO/Al2O3 methanol synthesis catalyst, either as a physically separate particle or as a component loaded onto a solid acid base. A variety of solid acids have been investigated, including alumina and several zeolites. Aguayo et al.9 tested a blend of methanol synthesis catalyst, and either alumina or NaHZSM-5 zeolite. They found, with the zeolite blend, that maintaining a slight steam vapor pressure at 260 C helped preserve catalyst activity, apparently due to the steam gasification of coke precursors. The alumina blend, however, did not respond favorably to the steam treatment because the alumina was not stable in the presence of steam. Fei et al.10 tested a series of acid catalysts based on HY zeolite, in which the hydrogen ion was exchanged with one of a number of metal ions. Prolonged testing at 245 C showed that two of these catalysts, the nickel and zirconium forms, gave good activity with minimal coke formation. The number and strength of acid sites on these catalysts were tested using ammonia temperature programmed desorption (TPD). Three acid strengths were found, a weak acid site with a TPD peak at about 220 C, a medium acid strength site with a TPD peak from about 280 to about 420 C, and a high acid strength site with a TPD peak at about 580 C. The two best catalysts, those with both high activity and little deactivation, had a high concentration of medium strength peaks and no measurable high strength peaks.

Methanol and Derivatives

One approach to improving temperature control in DME synthesis is to use a slurry bubble column reactor. The solid catalyst is suspended in a liquid oil bath, while feed gas bubbles into the bottom of the reactor. The oil, which is chemically inert, serves to transfer heat from the catalyst to cooling tubes. Lee and Gogate11 carried out a laboratory investigation of this reaction system using a stirred autoclave reactor. They chose Witco 40, a naphthenic oil in the light vacuum gas oil boiling range, as their process oil. Later, Air Products Company12 scaled this process up to a nominal 10 ton/day rate at their LaPorte, Texas pilot plant. Earlier, they tested a slurry bubble column reactor to synthesize methanol. The catalyst in this reactor was changed to synthesize DME. JFE Holdings, starting in 2002, built a 100 ton/day demonstration plant in Kurshiro, Japan, using a similar process.13 The slurry bubble reactor is designed to operate at 5 MPa and 260 C. Slurry bubble column reactors were reviewed by Kantarci et al.,14 primarily for application to biochemical reactions. These reactors contain turbulent mixtures of solids, liquid, and gas. The complex and coupled fluid mechanics, heat transfer, mass transfer, and chemical reactions in these reactors do not readily lend themselves to straightforward reactor design and control.

DME TO HYDROCARBONS: EXXONMOBIL MTG PROCESS In the early 1970s, researchers at Mobil Oil, now part of ExxonMobil, discovered that they could limit the molecular weight of hydrocarbons formed by the dehydration of DME (R-12.7) by using zeolite catalysts with well-defined pore sizes. In particular, hydrocarbons formed on the newly discovered ZSM-5 zeolite nearly all fit in the gasoline boiling range. This formed the basis of the ExxonMobil MTG, Methanol To Gasoline, process.15e17 The Mobil researchers obtained a large number of closely related patents, so the preferred catalyst was not clearly identified. The original patents expired, but the process continues to be protected by trade secrets. Haldor Topsoe offers a similar process, called TIGAS (Topsoe Integrated Gasoline Synthesis).18 The dehydration of methanol to gasoline is a multistep process. Measurements15 of product composition versus reactor space time reveal the following sequence of steps: CH3 OH/CH3 OCH3 ðþH2 OÞ/C2 to C5 alkenesðþH2 OÞ/C6 þ hydrocarbons R-12.8 The first commercial MTG plant started up in1985 in Motunui, New Zealand, with a capacity of 2,200 tons/day of gasoline.15 Methanol was produced using natural gas. The methanol was heated to 300 to 320 C, and fed to a reactor packed with an alumina catalyst for conversion of the methanol to DME and water (this is a two pot DME synthesis). Effluent from the DME reactor at about 400 to 420 C was mixed with recycle gas and fed to the top of the MTG reactors, packed with ZSM-5 catalyst. As DME is

365

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Methanol and Derivatives

converted to hydrocarbons, the catalyst is deactivated by coke. The reaction occurs within a narrow band of the catalyst bed, since the catalyst above this band is deactivated. The Motunui plant used five reactors, with four on line. In the fifth reactor, the catalyst was regenerated by burning off the coke with a mixture of oxygen and nitrogen. A slower, permanent catalyst deactivation was observed due to dealumination of the catalyst by steam. The reactor effluent was cooled to produce a three phase mixture including gas, the product hydrocarbons, and water. Hydrocarbon processing16 is shown in Figure 12.9. Hydrocarbon liquids from the phase separator are fed to a sequence of three distillation columns. The first column removes C2 and lower boiling compounds. The second column produces an LPG overhead, primarily a mixture of C3 and C4 hydrocarbons. The boiling range of this LPG product can be controlled to meet gasoline volatility specifications. Gasoline from the bottom of the second column is fed to a third column, where it is split into light gasoline and heavy gasoline streams. This is because the heavy gasoline contains substantial quantities of durene (1,2,4,5-tetramethyl benzene) (Figure 12.10), which, due to its symmetry, has a 79 C melting point. The heavy gasoline is fed to the HGT (heavy gasoline treatment) reactor where it is isomerized. Isomers of durene have much lower melting points. For example, 1,2,3,5-tetramethyl benzene, which lacks the symmetry of durene, has a 23.7 C melting point. Some molecular weight reduction

DME + H2O

purge

C2-

LPG

light gasoline

N2, CO2 recycle gas offline MTG reactor burning coke off catalyst

heavy gasoline

online MTG reactor

LPG

compressor

gas

hydrocarbon liquids

N2, O2

HGT reactor

water

Figure 12.9 ExxonMobil MTG (methanol to gasoline) process.15e17 The feedstock is a mixture of dimethyl ether (DME) and water produced from methanol.

Methanol and Derivatives

H3C

CH3 C

C

C

C C

H3C

C CH3

Figure 12.10 Durene (1,2,4,5-tetramethyl benzene) is present in substantial concentrations in the MTG reactor efﬂuent. Because of its high melting point, 79 C , the heavy gasoline must be isomerized in the HTG reactor.

occurs in the HTG reactor, so the effluent is distilled in order to meet gasoline volatility specifications. This produces a second LPG by-product. Table 12.1 shows hydrocarbon product yields and gasoline composition reported for the Montunui plant.15 Gasoline volatility specifications vary according to the season and the local climate. Allowable volatilities are higher during the winter than during the summer. Gasoline volatility is controlled primarily by adjusting the C4 content of the gasoline, so the gasoline yield should be somewhat higher than the 79.9 wt.% shown in Table 12.1. This is a remarkably high selectivity for gasoline. Table 12.2 compares MTG gasoline properties to typical US summer and winter gasoline blends.17 The MTG gasoline fits well within the normal spread of gasoline properties. The MTG reactor effluent requires minimal processing to produce a drop-in replacement for gasoline produced from conventional crude oil. In comparison, gasoline can also be made from syngas via Fischer-Tropsch synthesis, but, as shown in the next chapter, this gasoline has a very low octane number and requires considerable processing to meet current specifications. The Motunui plant was shut down when a drop in crude oil prices made the gasoline less valuable than the methanol produced by the plant. ExxonMobil claims to have made numerous improvements to the process since the Montunui plant was shut down. DKRW plans to build an MTG plant near Medicine Bow, Wyoming, which will produce 20,000 bbl/day of gasoline. The feedstock for this plant will be bituminous coal Table 12.1 Hydrocarbon yields and gasoline composition for the ExxonMobil MTG plant at Monutnui, New Zealand.15 Hydrocarbon product wt.% Gasoline composition

Light gas Propane Propene Isobutane n-Butane Butenes C5þ gasoline

1.4 5.5 0.2 8.6 3.3 1.1 79.9

Branched alkanes Branched alkenes Naphthenes (cycloalkanes) Aromatics

wt.%

53 12 7 28

367

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Methanol and Derivatives

Table 12.2 Comparison of ExxonMobil MTG gasoline properties to typical 2005 US gasolines.17 MTG gasoline Summer gasoline Winter gasoline

Oxygen, wt.% API gravity Aromatics, vol.% Olefins (alkenes), vol.% Reid vapor pressure, psi 50% boiling point, F 90% boiling point, F Sulfur, ppm Benzene, vol. %

0 61.8 26.5 12.6 9 201 320 0 0.3

0.95 58.4 27.7 12 8.3 211.1 330.7 106 1.21

1.08 61.9 24.7 11.6 12.12 199.9 324.1 97 1.15

from an adjacent, dedicated, underground coal mine operated by Arch Coal. The plant will use a GE gasifier and a Davy methanol plant. Due to limited water availability, the plant will make extensive use of air cooling instead of the more common evaporative cooling. DKRW originally intended to use a Fischer-Tropsch process, and then changed to the ExxonMobil MTG process.

METHANOL TO HYDROCARBONS: UOP/HYDRO MTO PROCESS The UOP/HYDRO MTO (methanol to olefins) process19 uses chemistry that is similar to that used in the ExxonMobil MTG process to produce olefins (alkenes) from methanol. Conventionally, olefins are made by pyrolyzing alkane feedstocks. In countries with natural gas production, the feedstock is generally natural gas liquids, which are C2þ hydrocarbons separated from natural gas. In countries that lack natural gas production, the feedstock is generally naphtha (unfinished gasoline). The feedstocks are mixed with steam to reduce coke deposition, and then heated above their pyrolysis temperatures in tubes inside a furnace. Ethylene, one of the main products, is use to make polyethylene, the most common type of polymer produced. Propylene is used to produce polypropylene, a major polymer type. Some polybutylene is made from butylene, but much of the butylene is used to produce isooctane via the alkylation process for gasoline. Butadiene is a major ingredient in synthetic rubber production. The UOP/HYDRO MTO process is shown in Figure 12.11. Methanol is fed to the bottom of a fluidized bed reactor, where methanol first dehydrates to form DME and then further dehydrates to form olefins. CH3 OHðgÞ4CH3 OCH3 ðgÞ þ H2 OðgÞ/C2 to C4 olefinsðgÞ þ H2 OðgÞ (R-12.9) This is essentially the same sequence of steps used in the ExxonMobil MTG process (R-12.8), except that (1) the methanol dehydration steps and the DME dehydration steps

Methanol and Derivatives

CO2 removal

CO2

methane

compressor

compressor H2

water removal

demethanizer water

dehydration

water

Acetylene saturator

CO2, N2

ethylene

MTO reactor

propylene

deethanizer depropanizer regenerator methanol

air

C4+ product

Figure 12.11 UOP/HYDRO MTO (methanol to oleﬁns) process.

are in the same reactor, and (2) the catalyst and reactor conditions are chosen such that most of the olefins do not oligomerize to form higher molecular weight hydrocarbons. Like the MTG process, coke forms on the catalyst during DME dehydration. A stream of catalyst is sent to the regenerator, another fluidized bed reactor, where the coke is burned off with air. A stream of regenerated catalyst flows from the regenerator to the MTO reactor. This combination of a primary fluidized bed reactor, with coke lay-down, and secondary fluidized bed reactor, with coke burn-off, is very similar to the fluidized bed catalytic cracking units commonly used in petroleum refining. This use of fluidized bed reactors requires a more physically robust catalyst (to prevent attrition) than that required for a packed bed reactor, such as the MTG reactors used in Montunui, New Zealand. Product gases from the MTO reactor flow to a water removal step. The water removal technique is not clearly specified in the process literature, but it is probably condensation. The gases then flow to a carbon dioxide removal step. Again, the technique is not specified, but one of the acid gas removal processes described in Chapter 6 would meet this purpose. The gas is compressed and sent to a dehydration step, which could be a glycol-based absorption process or a solid desiccant, such as silica gel. The deethanizer splits the gases into a C2 and lighter overhead and a C3 and heavier bottoms. The C2 and lighter stream is catalytically hydrogenated to convert acetylene to ethylene: o C2H2ðgÞ þ H2ðgÞ/C2H4ðgÞ DHrxn ¼ 174:3 kJ=gmole

R-12.10

369

370

Methanol and Derivatives

The treated gases are split in the demethanizer into an ethylene bottoms product stream and an overhead stream containing lighter gases. The bottoms stream from the deethanizer is fed to the depropanizer, where it is split into a propylene overhead product stream and a bottoms heavy product stream. Reactor conditions may be adjusted to emphasize either ethylene or propylene production.

METHANOL AND DME TO HYDROCARBONS: FUTURE PROSPECTS Dehydration of methanol to DME, and subsequent dehydration to hydrocarbons appears to be a versatile approach for the conversion of synthesis gas to hydrocarbons. Additional hydrocarbon products, such as jet fuel or diesel, are possible with the proper selection of catalyst and reaction conditions. Compared to Fischer-Tropsch synthesis, the methanol/ DME route is relatively simple and gives high yields of the desired hydrocarbon products.

REFERENCES 1. Bridger GW, Spencer MS. Methanol synthesis. In: Twigg MV, ed. Catalyst Handbook. 2nd edn. Wolfe Publishing; 1989. 2. Natta G. Synthesis of Methanol, in Catalysis, Vol. III, Hydrogenation and Dehydrogenation, Emmett PH. ed. 1955;349e411. 3. Graaf GH, Stamhuis EJ, Beenackers AACM. Kinetics of low pressure methanol synthesis. Chem Eng Sci. 1988;43:3185-3195. 4. Graaf GH, Scholtens H, Stamhuis EJ, Beenackers AACM. Intra-particle diffusion limitations in lowpressure methanol synthesis. Chem Eng Sci. 1990;45:773-783. 5. Du Pont. Toxicity summary for dimethyl ether (DME); DymelR A Propellant www.aboutdme.org/ EFIClient/files/ccLibraryFiles/Filename/000000000640/DME_Technical_-_Toxicity_Summary.pdf; 1987. 6. The Catalyst Group, Global Dimethyl Ether Emerging Markets. April 20, 2007, www.aboutdme.org/ EFIClient/files/ccLibraryFiles/Filename/000000000424/2007_DME_Seminar_9_Payne_CatalystGroup. pdf. 7. Holm-Larsen H. Synthesis and New Applications of DME e A Review of Alternatives, May 30, 2007, www.aboutdme.org/EFIClient/files/ccLibraryFiles/Filename/000000000737/2007_NGCS8_HolmLarsen_HaldorTopsoe.pdf. 8. Ekstrand J, Laegsgaard Jorgensen S. Cost Effective Topsoe DME Production Technology, Gasochem 2007, <www.topsoe.com/business_areas/gasification_based/Processes/~/media/PDF%20files/DME/ Topsoe%20DME%20%20%20Gasochem%202007.ashx>. 9. Aguayo AT, Eren˜a J, Sierra I, Olazar M, Bilbao J. Deactivation and regeneration of hybrid catalysts in the single-step synthesis of dimethyl ether and CO2. Catal Today. 2005;106:265-270. 10. Fei J, Hou Z, Zhu B, Lou H, Zheng X. Synthesis of dimethyl ether (DME) on modified HY zeolite and modified HY zeolite-supported Cu-Mn-Zn catalysts. Appl Catal A Gen. 2006;304:49-54. 11. Lee S, Gogate M. Dimethyl ether synthesis process, Electric Power Research Institute TR-100246, Project. 1992;317-6. 12. Air Products and Chemicals, Inc. Liquid phase dimethyl ether demonstration in the LaPorte Alternative Fuels Development Unit, Topical Report, U.S. Dept. of Energy Agreements DE-FC2292PC90543 and DE-FC22-95PC93052; 2001. 13 de Meister du Bourg H. Future Prospective of DME, 23rd World Gas Conference, Amsterdam; 2006. 14. Kantarci N, Borak F, Ulgen KO. Review: Bubble column reactors. Proc Biochem. 2005;40:2263-2283.

Methanol and Derivatives

15. New Zealand Institute of Chemistry, The production of methanol and gasoline, in Chemical Processes in New Zealand, Energy chapter; 1998. nzic.org.nz/ChemProcesses/energy/7D.pdf. 16. Zhao X, McGihon RD, Tabak SA. Coal to clean gasoline. Hydrocarbon Engr. 2008;13(9):39-49. 17. ExxonMobil Research and Engineering, Methanol to Gasoline (MTG) production of clean gasoline from coal. 18. Haldor Topsoe, Gasoline-Tigas www.topsoe.com/business_areas/gasification_based/Processes/Gasoline_ TIGAS.aspx. 19. UOP, UOP/HYDRO MTO Process methanol to olefins conversion, <www.uop.com/objects/26% 20mto%20process.pdf>.

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CHAPTER

13

Substitute Natural Gas and Fischer-Tropsch Synthesis Contents Overview Substitute Natural Gas SNG Processes Competition Between Coal Gasiﬁcation and Natural Gas SNG as a Carrier of Coal Energy SNG Versus Synthetic Liquid Hydrocarbons Fischer-Tropsch History Fischer-Tropsch Chemistry FT reactor design Reﬁning Fischer-Tropsch Fluids Diesel Production Naphtha and Gasoline Lubricating Base Oils Outlook for FT Fluid Reﬁning FT Economics References

373 373 375 377 378 378 379 379 381 385 385 386 387 388 389 390

OVERVIEW This chapter covers the direct conversion of syngas to hydrocarbons, either substitute natural gas, or higher hydrocarbons via Fischer-Tropsch synthesis. In the prior chapter, an indirect route via methanol and dimethyl ether synthesis was considered. Fischer-Tropsch synthesis produces a synthetic crude oil, and the unique properties of this oil require special considerations during refining to produce petroleum products.

SUBSTITUTE NATURAL GAS Substitute natural gas (SNG), also known as synthetic natural gas, is methane. Methane is formed during gasification, and Chapter 3 briefly described gasification processes designed to produce methane as their primary product.1,2 The usual approach, however, is the catalytic hydrogenation of carbon monoxide: COðgÞ þ 3H2 ðgÞ/CH4 ðgÞ þ H2 OðgÞ

DH rxn ¼ 206:2 kJ=gmole

Coal Gasification and Its Applications. ISBN B978-0-8155-2049-8.10013-0, doi:10.1016/B978-0-8155-2049-8.10013-0

R-13.1

Ó 2011 Elsevier Inc. All rights reserved.

373

Substitute Natural Gas and Fischer-Tropsch Synthesis

The syngas will generally contain carbon dioxide, and this, too, can be hydrogenated to produce methane: CO2 ðgÞ þ 4H2 ðgÞ/CH4 þ H2 OðgÞ DH rxn ¼ 165:0 kJ=gmole

R-13.2

SNG catalysts, usually nickel, also have significant water gas shift activity: COðgÞ þ H2 OðgÞ4CO2 ðgÞ þ H2 ðgÞ

DH rxn ¼ 41:21 kJ=gmole

R-13.3

(a)

Equilbirum Conversion of CO,

Figure 13.1a shows the equilibrium conversion of carbon monoxide to methane via R-13.1 as a function of temperature and pressure. Conversion falls with increasing temperature. Reaction R-13.1 shows that four moles of gas react to form two moles of gas, so increasing reaction pressure pushes the reaction to the right. In all of the equilibria calculated for Figure 13.1a, the conversion of carbon monoxide was 99.2% or higher.

100 99 98 97 96 95 94 93 92 91 90 260

5 MPa 2 MPa 1 MPa

280

300

320

340

360

380

400

420

440

460

480

500

Temperature, C

(b)

Equilibrium Conversion of CO,

374

9 8

1 MPa

7 2 MPa

6 5

5 MPa

4 3 2 1 0 260

280

300

320

340

360

380

400

420

440

460

480

500

Temperature, C

Figure 13.1 (a) Equilibrium conversion of carbon monoxide to methane via R-13.1. Most of the drop in conversion with increasing temperature is due to conversion of carbon monoxide to carbon dioxide via the water gas shift reaction (R-13.3). A stoichiometric 3:1 hydrogen to carbon monoxide feed ratio was used in these calculations. (b) Equilibrium conversion of carbon monoxide to carbon dioxide via the water gas shift reaction (R-13.3) during the conversion of carbon monoxide to methane via R-13.1.

Substitute Natural Gas and Fischer-Tropsch Synthesis

The reason for the decline in conversion of carbon monoxide to methane with increasing temperature is the competing water gas shift reaction, R-13.3. Figure 13.1b shows how the equilibrium carbon dioxide concentration increases as a result of the water gas shift reaction with increasing reaction temperature. The conversion of hydrogen was as high as 99.6% at the minimum temperature, 260 C, and the maximum pressure, 5 MPa, and as low as 87.9% at the maximum temperature, 500 C, and the minimum pressure, 1 MPa. Figure 13.2 shows the equilibrium conversion of carbon dioxide to methane via R-13.2. Methane yields are a little lower for carbon dioxide compared to carbon monoxide at the same reaction conditions. For carbon dioxide, a lower yield of byproducts are obtained. Less than 0.9% of the carbon dioxide is converted to carbon monoxide via the reverse water gas shift reaction, R-13.3.

SNG PROCESSES The Great Plains Synfuels plant,3 in Beulah, North Dakota, started operation in 1984. Lignite from an adjacent mine is crushed and fed to Lurgi gasifiers. Since these gasifiers cannot process fine coal, fines are sent to an adjacent pulverized coal combustion power plant. Syngas is cleaned in a Rectisol unit, also a Lurgi process. Carbon dioxide from the Rectisol unit is pipelined to Weyburn, Saskatchewan, where it is used for enhanced oil recovery. Cleaned syngas from the Rectisol unit is sent to a Lurgi methanation unit, where it is converted to SNG. The catalyst in the methanation unit is gradually poisoned by sulfur, so the catalyst lifetime is dictated by operation of the Rectisol unit, which removes the sulfur.

Equlibrium Conversion of CO2,

100

95

5 MPa

2 MPa

90 1 MPa

85 260

280

300

320

340

360

380

400

420

440

460

480

500

Temperature, C

Figure 13.2 Equilibrium conversion of carbon dioxide to methane via R-13.2. The feed for these calculations was a stoichiometric 4:1 hydrogen to carbon dioxide ratio.

375

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Substitute Natural Gas and Fischer-Tropsch Synthesis

H2, CO, CO2

reactor 1

reactor 2

reactor 3

recycle gas

steam cooler

water

steam water

boiler feed water preheater cooler methane

recycle compressor water

Figure 13.3 Haldor Topsoe TREMPÔ (Topsoe Recycle Energy-efﬁcient Methanation Process) for substitute natural gas.

Figure 13.3 shows a methanation unit design by Haldor Topsoe, called TREMPÔ (Topsoe Recycle Energy-efficient Methanation Process).4 A nearly stoichiometric blend of hydrogen, carbon monoxide, and carbon dioxide is mixed with recycle gas and fed to the first reactor. The effluent from this reactor is cooled, and then split into a recycle stream, and the feed is then sent to the second reactor. The recycle stream is further cooled in two exchangers. This gas is then fed to the recycle compressor. The recycle gas dilutes the feed gas, and reduces the temperature rise across the first reactor. A similar recycle gas cooling scheme is used in the Lurgi methanation process,5 in which the gas enters the first reactor at 300 C and leaves the first reactor at 450 C. The Lurgi methanator operates at about 1.8 MPa. Unlike ammonia or methanol synthesis, there is no separation of unreacted feed gas and product. With SNG, this would require an expensive cryogenic distillation process. Instead, the TREMPÔ SNG process is designed to achieve high single pass conversion. This is done by a series of three reactors, each reactor operating at a lower temperature than the previous reactor. As can be seen in Figures 13.1a,b, and 13.2, lower operating temperatures favor the formation of methane, and reduce the formation of carbon dioxide byproduct. Table 13.1 shows typical gas compositions and heating values for SNG produced by the TREMPÔ process.4

Substitute Natural Gas and Fischer-Tropsch Synthesis

Table 13.1 Typical gas composition and heating value for substitute natural gas produce by the TREMPÔ process.4 Component Mole %

Methane Carbon dioxide Hydrogen Carbon monoxide Nitrogen þ argon Higher heating value, KJ/Nm3

94e98 0.2e2 0.5e2 <100 ppm 2e3 37,380e38,370

COMPETITION BETWEEN COAL GASIFICATION AND NATURAL GAS Coal gasification and natural gas have a long history of commercial competition. The first coal gasifiers produced town gas for household energy use. These gasifiers were shut down when natural gas and electric power became widely available. The first ammonia and methanol plants were designed to use syngas from coal gasifiers. Now, most ammonia and methanol is made from natural gas-derived syngas. In the late 1960s through the early 1980s, there was widespread concern that North America was about to exhaust its natural gas reserves. This led to the construction of the Great Plains Synfuels plant. Natural gas reserves were not near exhaustion, and in the mid 1980s the price of natural gas fell. The Great Plains Synfuels plant was uneconomic, and the US government intervened to keep it operating. In the early 2000s rapidly rising natural gas prices had a number of companies considering ammonia and methanol production from coal, but natural gas prices fell again. When this book was written, the outlook was for abundant natural gas supplies and little likelihood of large price increases. Two developments are responsible for this situation. First, consider that natural gas markets traditionally are regional, rather than global. Regional producers are connected to regional consumers via a pipeline network. The USA imported most of its crude oil in 2009, but the USA produced 92% of the natural gas that it consumed that year. The imported gas was almost entirely from Canada, and shipped to the USA via pipeline. A similar regional market exists in Europe, with producers in the North Sea, Russia, and northern Africa. In some regions, such as east Asia, the potential consumption of natural gas is greater than the regional supply. In other regions, such as the Persian Gulf, regional natural gas reserves are large compared to regional consumption. One of the two major developments that affected the natural gas outlook recently is liquefied natural gas (LNG), which is converting the natural gas market into a global market. In regions with an abundance of natural gas, such as the Persian Gulf, natural gas is liquefied and loaded onto transoceanic tankers. It is then shipped to regions where local demand exceeds supply, such as east Asia, where it is vaporized and put into the

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local pipeline network. No great technological breakthrough was necessary to develop this market. Instead, sustained price differentials between supplying regions and consuming regions made LNG profitable. The second major development is a combination of two technical advances. The first is directional drilling, which allows the driller to follow the length of natural gas-bearing reservoir rock, rather than merely drilling a vertical hole through it. The second technical development is hydraulic fracturing, which fractures reservoir rock around the well bore. These two developments allow drillers to create large drainage areas in each well. Now, natural gas production from low permeability rocks, especially shale, is economically feasible. Huge volumes of shale in North America have been recently re-classified as natural gas reserves. Shale gas production, now common in North America, is expected to spread to other continents. So where does coal gasification fit in? Well, for many applications, such as ammonia and methanol production, the abundance of natural gas put gasification plans on hold. Uses of natural gas are expanding, such as an increase in natural gas for electric power production, as opposed to coal combustion, to reduce greenhouse gas production, and the potential use of natural gas as a transportation fuel to replace scarce petroleum. Of course, natural gas will not last forever, and SNG is a valuable back-up supply of methane when natural gas supplies diminish.

SNG AS A CARRIER OF COAL ENERGY Currently, most coal is burned to produce electric power. Coal is hauled over land in trains, and a large fraction of a power plant’s delivered coal cost is transportation cost. Natural gas, on the other hand, may be transported long distances at relatively low cost in pipelines. Compared to coal, natural gas combustion produces much less carbon dioxide per unit of energy. Rather than shipping coal in a train, coal could be converted to SNG in a plant near the coal mine. By-product carbon dioxide could be locally sequestered. The SNG could then be shipped via pipeline to power plants and other consumers that might otherwise burn coal. This approach will lower transportation costs and greenhouse gas emissions.

SNG VERSUS SYNTHETIC LIQUID HYDROCARBONS Traditionally, liquid petroleum products were more expensive than natural gas on an equivalent energy basis. This is why burning oil to produce heat or electricity is generally limited to regions without natural gas supplies. Lately, the spread between oil and natural gas prices has widened considerably. If current projections about crude oil supply and demand come true, that gap will widen further.

Substitute Natural Gas and Fischer-Tropsch Synthesis

As shown in the previous chapter, liquid hydrocarbons can be made from syngas via methanol and dimethyl ether. This chapter discusses a second, better-known approach which is the Fischer-Tropsch synthesis. If a decision between converting syngas to SNG or liquid fuels needs to be made, one should remember that the liquids are worth much more.

FISCHER-TROPSCH HISTORY During the 1920s, Franz Fischer and Hans Tropsch, working at the Kaiser Wilhem Institute in Berlin, developed a process for the direct conversion of syngas to hydrocarbons. This is known today as the Fischer-Tropsch (FT) process. During World War II, Germany, cut off from crude oil suppliers, produced synthetic oil using the FT process and the Bergius process, which is a direct coal hydrogenation process (see Chapter 2). Today, the Fischer Tropsch Archive5 maintains an electronic collection of early FT process documents. Sasol, using Lurgi technology, started production of FT fluids in 1955 from locally mined coal in Sasolburg, South Africa.6 A second FT plant in Secunda, South Africa, started operations in 1980. Much of the FT process technology literature is based on Sasol’s South African experience. Shell started their SMDS (Shell middle distillate synthesis) plant in Bintulu, Malaysia in 1993. This plant uses natural gas as the feedstock, and the FT fluids are hydrocracked. The Oryx GTL (gas to liquids) plant, a joint venture between Qatar Petroleum and Sasol, started operations in 2006. The Pearl GTL plant, a joint venture between Qatar Petroleum and Shell, is expected to be the world’s largest gas to liquids plant when it reaches its full 140,000 bbl/day liquids production in 2011.

FISCHER-TROPSCH CHEMISTRY Fischer-Tropsch synthesis is a polymerization reaction. An alkane chain, absorbed on a catalyst site ()), reacts with carbon monoxide and hydrogen to add a methyl unit (CH2): H-ðCH2 Þn- þCO þ 2H2 /H-ðCH2 Þn CH2- þ H2 O

R-13.4

The heat of reaction per methyl group, estimated from the gas phase heat of formation for dodecane (C12H26), is 155.5 kJ/mole. This heat of reaction is about 20% of the heat of combustion of the feed gases. Iron and cobalt catalysts are typically used, although other metals, such as ruthenium and nickel, are also active. Iron has significant water gas shift activity (R-13.3), and typical FT synthesis conditions favor the right side of the water gas shift reaction. Combining R-13.3 with R-13.4 for an iron catalyst gives: H-ðCH2 Þn- þ2CO þ H2 /H-ðCH2 Þn CH2- þ CO2

R-13.5

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The heat of reaction per methyl group, again based on dodecane, is 196.7 kJ/mole. This is somewhat higher than R-13.4 due to the exothermic contribution of the water gas shift reaction. This heat of reaction is about 24% of the heat of combustion of the feed gases. An examination of R-13.4 and R-13.5 shows that, for a cobalt catalyst, the ideal feedstock has a 1:2 carbon monoxide to hydrogen ratio, while, for an iron catalyst, the ideal feedstock has a 2:1 carbon monoxide to hydrogen ratio. With a typical syngas from a coal gasifier, a water gas shift reactor is required for a cobalt catalyst, but not required for an iron catalyst. The water gas shift reactor, could, however, shift part of the heat removal from an iron FT synthesis reactor to the water gas shift reactor, and, as we will see later, this could improve operations. The absorbed alkane chain may desorb, halting growth. The usual desorption product is a normal alkane, but alcohols and olefins also form. This means that the final product has a range of molecular weights, ranging from methane to very high molecular weight alkanes. The choice of catalyst and operating conditions determine the distribution of molecular weights. Anderson7 adapted a model developed for other types of polymerization to FT synthesis. The result is known as the Anderson-Schultz-Flory (ASF) distribution. A single adjustable parameter, a, describes the probability that an alkane absorbed on the catalyst, H-(CH2)n-), will add a methyl unit to form H-(CH2)nþ1-). Conversely, (1a) is the probability that the absorbed alkane will desorb, halting growth. The final weight fraction of each alkane compound (using the approximation that only alkanes are formed) is: W ¼ nan-1 ð1-aÞ2 Eqn. 13.1 n

Variations on the ASF distribution equations are used to improve the fit to experimental measurements of molecular weight distributions. For example, olefins appear to reabsorb and re-initiate growth, so high molecular weight alkanes are a little more abundant than predicted by Eqn. 13.1. Except for a relatively small quantity of hydrocarbon solvents, petroleum products are not sold as pure compounds. Instead, petroleum products are sold as hydrocarbon mixtures that are primarily distinguished by their boiling point ranges. There can be considerable overlap between adjacent products. For example, some of the higher molecular weight compounds in gasoline can also be some of the lower molecular compounds in diesel. Still, it is useful to approximate the FT fluid distillation product yield using molecular weight ranges. The approximate distillation yield for an FT fluid versus the chain growth probability factor, a, is shown in Figure 13.4. The catalyst, reactor design, and operating conditions are all optimized to maximize the value of the hydrocarbon products. In a typical petroleum refinery, the refinery may be operated to maximize the yield of naphtha, which is further processed to make

Substitute Natural Gas and Fischer-Tropsch Synthesis 100

Fuel gas (C1 + C2)

90

Distillation Yield, wt.

80

LPG (C3 + C4)

70 60 Naphtha(C5 - C11)

50 40

Diesel or Jet Fuel (C12 - C18)

30 20

Wax (C19+)

10 0 0.75

0.77

0.79

0.81

0.83

0.85

0.87

0.89

0.91

0.93

0.95

Figure13.4 Product distillation yield for a Fischer-Tropsch liquid as a function of a, the chain growth probability factor. Carbon numbers are used to estimate distillation boiling ranges. Yield versus carbon number was estimated using Eqn. 13.1.

gasoline. FT naphthas, on the other hand, tend to have very low octane numbers, while FT diesels tend to have very high cetane numbers. Lamprecht8 reports a 62 cetane number for an FT diesel, versus a 40 minimum cetane specification9 for US diesel. Both the excellent FT diesel cetane number and the poor FT naphtha octane number are for the same reason, a high normal alkane content. Therefore, FT plants tend to maximize diesel production. The least valuable fraction is the fuel gas. In a typical oil refinery, this gas is burned to provide heat for refinery operations. The next least valuable fraction is LPG. This may be recovered and sold, but the price is substantially less than for gasoline, jet fuel, or naphtha. For example, in 2009, the average US price for propane was $1.236/ gallon, versus $1.888, $1.704, and $1.833 per pre-tax gallon for gasoline, jet fuel, and diesel, respectively.10 Many refineries have dimerization processes that convert LPG to naphtha. Demand for paraffin wax is small, so the usual approach is to hydrocrack FT wax to produce lower molecular weight hydrocarbons. To maximize production of diesel, the FT reactor is designed to operate with a high chain growth probability factor, a, and then the wax is hydrocracked.

FT REACTOR DESIGN Lu11 reviewed a number of FT experimental studies; and developed approximate correlations between the chain growth probability factor, a, and reactor temperature and pressure. The most prominent trend is that a tends to decline with increasing

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temperature. This means that the operating temperature is a compromise between good reaction rates, they increase with increasing temperature; and high molecular weight distributions, which increase with decreasing temperature. A further challenge is that FT synthesis is highly exothermic, so the reactor must be designed to remove a large quantity of heat to maintain the reactor temperature within an optimum operating temperature range. Figure 13.5 shows the Sasol Synthol circulating fluid bed reactor. Fresh and recycled gas are mixed with a hot catalyst stream and fed to a vessel containing cooling bundles. The catalyst/gas mixture is then fed to a disengagement vessel where product gas is taken off the top, through cyclones to remove fine catalyst solids. The catalyst flows through a standpipe on the bottom of the disengagement vessel. The first Sasol Advanced Synthol reactor12 started production in Sasolburg in 1989. This reactor, shown Figure 13.6, is a bubbling fluidized bed reactor with heat transfer tubes in the catalyst bed. Compared to the original Synthol reactor, the Advanced Synthol reactor is more compact for a similar production capacity. The Synthol reactors operate at a fairly high temperature, 330 to 350 C. Sie and Krishna13 analyzed fluidized bed FT reactors, such as the Synthol reactors. They noted that because of the small catalyst diameter, 100 mm, there are no intraparticle diffusional

catalyst + gas product gas cyclones

cooling bundles catalyst

Fresh gas + recycle

catalyst + gas

Figure 13.5 The Sasol Synthol circulating ﬂuid bed reactor.

Substitute Natural Gas and Fischer-Tropsch Synthesis

product gas steam

water syngas

Figure 13.6 The Sasol Advanced Synthol reactor is a bubbling ﬂuidized bed reactor with heat transfer tubes in the catalyst bed.

restrictions. The reactor temperature and chain growth factor, a, are chosen to prevent the condensation of liquid products in the catalyst bed. The Synthol reactor, shown in Figure 13.5, was one of two original reactor types used by Sasol. The other type is the ARGE reactor, shown in Figure 13.7. This reactor is built like a shell-and-tube heat exchanger, with catalyst packed in the tubes and coolant circulating through the shell side. Both gaseous and liquid hydrocarbons are formed, so syngas is fed to the top of the reactor and liquid wax trickles down through the packed

feed gas

steam

water product gas

wax

Figure 13.7 The Sasol ARGE reactor. Catalyst is packed in the tubes, and coolant circulates on the shell side.

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Table 13.2 Fischer-Tropsch product distribution for iron-based catalysts.14 Component, % Low temperature (220e250 C) High temperature (330e350 C)

CH4 C2-C4 olefins C2-C4 alkanes Naphtha Distillate Oils and waxes Oxygenates

4 4 4 18 19 48 3

7 24 6 36 12 9 6

catalyst tube. There are two product ports on the bottom of the reactor, one for liquid wax and the other for gas. The ARGE reactor is operated at about 220 to 250 C, a much lower temperature than the Synthol reactor. This results in a higher chain growth factor, a, leading to a heavier product, as shown14 in Table 13.2 for iron-based catalysts. Cobalt catalysts tend to give higher a values,15 which lead to higher molecular weight products. The Shell SMDS reactor in Bintulu, Malaysia uses a packed tube configuration similar to the ARGE reactor. More recently, slurry bubble reactors have been used for low temperature FT synthesis. The reaction media is a three phase churning mixture of solid catalyst, liquid wax product, and feed and product gas. Heat exchange tubes in the mixture remove the heat of reaction. Catalyst particle diameters14 are between 20 and 200 mm. The lower diameter limit is set by the difficulty of removing smaller solids from the product wax, and the upper diameter limit is set by the need to keep the catalyst particle suspended in the reaction media. There have been reports of operational difficulties with slurry bubble reactors. Figure 13.8 shows a microchannel reactor.16 This reactor is created by etching parallel grooves in flat metal plates. The plates are stacked to form flow channels bounded by

Figure 13.8 A pilot scale microchannel reactor for Fischer-Tropsch synthesis built by Velocys, Inc.

Substitute Natural Gas and Fischer-Tropsch Synthesis

adjacent plates. Alternating layers of flow channels are filled with catalyst. The channels between the catalyst channels contain a heat transfer fluid, flowing at right angles with respect to the reactant flow. Conceptually, this reactor is much like the packed tube ARGE reactor. The microchannel reactor, however, has a much higher heat transfer area/catalyst volume ratio and a much smaller flow channel cross section. Both channels allow for a much better control of the reaction temperature range.

REFINING FISCHER-TROPSCH FLUIDS FT fluid can be thought of as a synthetic crude oil. As such, it is an unusual crude oil that requires special refining considerations. A typical crude oil contains normal alkanes, branched and cyclic alkanes, aromatics, sulfur compounds, and nitrogen compounds. Unless the crude oil is produced using thermal techniques, crude oil contains essentially no olefins. Oxygen compounds are present in crude oil, but usually at low levels. FT fluid, on the other hand, consists primarily of normal alkanes, with lesser but significant levels of olefins and alcohols. FT fluid contains very few branched alkanes, and essentially no cyclic compounds, sulfur or nitrogen compounds. We17,18 modeled FT refineries using a pure component approach and the Aspen Plus process simulation software. Three scenarios were considered in this study; a refining scheme that maximized diesel production, another, very similar scenario that maximized jet fuel production, and a third refining scheme to produce lubricating base oils.

Diesel Production FT diesel fuel, due to its high alkane content, has very high cetane numbers. Blends of FT and conventional diesel are sold as premium products in Europe. Normal alkanes, however, have high melting points compared to branched or cyclic hydrocarbons. Normal hexadecane, for example, has a melting point of þ18 C, versus 23 C for its branched isomer, 3-methyl-pentadecane. An FT diesel produced simply by distilling an FT fluid would have poor low temperature properties. This may be one of the reasons why FT diesel is sold in Europe as a blend with conventional diesel, rather than as a pure product. Most FT diesel, however, is not simply an FT fluid distillation product, but a blend of this distillation product and products of hydrocracked FT wax. Sie19 studied the hydrocracking of normal alkanes, and presented evidence for a cyclopropane intermediate mechanism. As a result of this mechanism, propane is the smallest fragment produced by a hydrocracking reaction. This reaction converts a normal paraffin molecule into two smaller molecules. One of these is also a normal alkane, while the other is a branched alkane. This shows a close relationship between hydrocracking and isomerization. As an approximation, any of the carbonecarbon bonds within the molecule,

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except for the two closest to the ends, have an equal probability of fracturing. Sie’s data actually show that the bonds towards the center of the molecule are slightly more likely to fracture than bonds near the end of the molecule. Leckel and co-workers20-23 published several papers describing Sasol FT wax hydrocracking pilot plant experiments. The FT wax is hydrocracked and then distilled. The distillation bottoms, which are unconverted wax, are recycled to extinction. Hydrocracking severity (conversion) can be controlled using a number of process parameters, including catalyst acidity, reactor temperature, reactor pressure, and reactor space velocity. Leckel showed that low temperature properties of diesel improve with increased hydrocracker severity, which is expected since hydrocracking creates branched alkanes. Increased hydrocracking also resulted in lower diesel yields, since a larger fraction of the hydrocracked product was naphtha or lighter products. Thus, there is a close relationship between diesel yield and low temperature properties. In our diesel model, we distilled an FT fluid produced by a cobalt catalyst with a high chain growth factor, a. Both the wax and the diesel fractions were hydrocracked. The diesel fraction did not need to be hydrocracked to meet the boiling point specifications. Rather, this was done to increase the branched alkane/normal alkane ratio in the final diesel product. Sie’s rules were used to model the hydrocracker; and the extent of hydrocracking was arbitrarily adjusted so that our simulation results fit Leckel’s experimental results. Although we were able to estimate the diesel composition, we did not find a reliable means of predicting low temperature properties based on that composition. In the USA, diesel cloud points are specified9 according to the time of year and the local climate. The cloud point is the temperature at which wax crystals become visible as the sample is cooled. Below the cloud point, these wax crystals plug fuel filters. Krishna et al.24 offer an approximate method of predicting pour points based on composition. Pour points are typically a few degrees lower than cloud points. Leckel’s results show that FT diesel can meet low temperature specifications, but we are not currently able to reliably predict diesel yields without experimental data.

Naphtha and Gasoline Naphtha processing proved to be challenging, due to its low octane numbers for FT naphtha. Brent et al.25 investigated FT naphtha upgrading techniques. They approached two naphtha reforming process providers, and were told by both that FT naphtha was not a good reforming feedstock. Naphtha reforming increases octane numbers by increasing aromatic content, primarily through dehydrogenation of cycloalkanes. Instead, they chose to sell FT naphtha as an olefin production feedstock. FT naphtha, compared to other naphthas, is an excellent olefin production feedstock because of its high alkane content. There are two problems with this approach. The first is that this would put FT

Substitute Natural Gas and Fischer-Tropsch Synthesis

naphtha in direct competition with natural gas liquids, which sell at substantially lower prices than gasoline. The second is that olefin plants designed to use naphtha feed are primarily located in countries that do not have significant natural gas production. In our model, we sought to convert FT naphtha to gasoline. To do this, we used two processes. The first is the TIP (total isomerization process) process,26 in which the octane number is increased by converting normal alkanes into branched alkanes. The isomerization reactor effluent is fed to a mole sieve bed, where normal alkanes are selectively adsorbed and recycled to the isomerization reactor. We used Sie’s hydrocracking rules to model the isomerization reactor, although catalysts with better isomerization/hydrocracking selectivity are available. The second process is an alkylation process that converts C4 compounds to isooctane. The problem with this refining approach was that octane improvement was accompanied by molecular weight reduction due to hydrocracking in the isomerization reactor, and we could not simultaneously meet both octane specifications and boiling point specifications. A more selective isomerization catalyst would improve the performance of this refining approach. Just as steam reforming of natural gas is a common method of producing syngas, LPG and naphtha can also be steam reformed to produce syngas.27 If a suitable outlet for FT naphtha or LPG is not found, then these streams can be recycled, thereby increasing the fractional yield of diesel or jet fuel. Our jet fuel model was very much like our diesel model, with the boiling point range adjusted to meet jet fuel standards. When jet fuel production is maximized, the diesel yield is insignificant.

Lubricating Base Oils We also examined a refinery model to maximize the production of lubricating base oils. This is the largest volume non-fuel petroleum market. A typical lubricating oil contains about 90% base oil and about 10% of an additive package. Branched alkanes, primarily in the C20 to C40 range, offer the best combination of volatility, viscosity, and low temperature performance for base oils. Traditionally, base oils with high branched alkane content are made using a paraffinic crude oil. A vacuum distillate cut is processed to increase the alkane content, either by solvent extraction of aromatics, or by converting aromatics to alkanes in a catalytic hydrogenation reactor. The oil is then dewaxed (high melting point normal alkanes are removed) by mixing the oil with a blend of methyl ethyl ketone and toluene, chilling the oil to crystallize wax, and then filtering the wax/ oil/solvent slurry. The solvent dewaxing process is the most expensive step in the traditional base oil production process. The traditional base oil process produces a complex and variable mixture of hydrocarbons. With increasingly stringent lubricating oil standards, producers found it difficult to meet standards with these traditional base oils. This led to the production of

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synthetic oils, such as polyalpha olefins, in which the hydrocarbon molecular structure can be more closely controlled to optimize oil properties. The cost of synthetic base oils is much higher than that of traditional base oils. In the Chevron Isodewaxing process28 normal alkanes are isomerized to branched alkanes. The pore size of the catalyst allows methyl branches to form, but the pores are too small to readily allow the formation of the cyclopropane intermediate compound that leads to hydrocracking. As a result, this catalyst has excellent isomerization/ hydrocracking selectivity. The Isodewaxing process was initially marketed as an inexpensive alternative to solvent dewaxing, but the isomerized waxes have physical properties that are superior to traditional base oils. When the feedstock is a paraffin wax, the base oil product is nearly as good as polyalpha olefins. This process offers a low cost route to high quality lubricants. Partly in response to this process, the American Petroleum Institute developed a classification system29 for base oils. Traditional base oils are either Class I or Class II, depending on their sulfur content. The higher quality base oils produced by wax isomerization are Class III. Poly alpha olefins are Class IV, and other types of synthetic base oils are Class V. In our lubricating base oil refinery model, the FT wax is isomerized using the Isodewaxing process. Detailed technical information on this process in the open literature is scarce, so we modeled the isomerization reactor based on pure component yield data from Claude et al.30-31 for a different catalyst, Pt/H-ZSM-22. This catalyst has a pore geometry that is similar to the Chevron catalyst. The yield data were for normal alkane feeds between C10 and C24. Yield data were extrapolated to higher molecular weights for this simulation. Three different viscosity grade oils were made. The base oil refinery model is attractive, in part, because base oils are worth considerably more than fuels. In March, 2010, US Gulf Coast Group III base oils sold for about $3.50/gallon,32 about twice the price of diesel.

Outlook for FT Fluid Reﬁning In our refinery models, we sought, to use conventional petroleum refining processes. The performance of the fuel refinery cases could be improved with the application of an isomerization catalyst with high isomerization/hydrocracking selectivity, such as the Chevron Isodewaxing catalyst used for base oil production. A combination of this catalyst and hydrocracking would give the refiner independent controls over molecular weight and low temperature properties, and this could greatly enhance diesel or jet fuel yields. It is not clear whether gasoline octane specifications can be met, but a higher octane is achievable if better isomerization selectivity is attained in the naphtha isomerization reactor. As a stand-alone process for liquid hydrocarbon production from syngas, FischerTropsch technology competes with methanol-based processes, such as the ExxonMobil

Substitute Natural Gas and Fischer-Tropsch Synthesis

MTG process discussed in the previous chapter. The MTG process is attractive because it produces specification grade gasoline in high yield with minimal processing. Fischer-Tropsch technology may be better suited as a companion technology for heavy oil refining. The trend in petroleum refining is away from light crude oils and towards heavy crudes. This produces much more aromatic products. High aromatic content in gasoline boosts octane numbers, but aromatic contents are limited due to toxicity concerns about aromatics, especially benzene. Aromatics tend to lower diesel cetane numbers. FT products are nearly free of aromatics, when blended with products from heavy crudes, the blends more closely resemble the mixture of alkanes and aromatics that are normally associated with refined products made from light crudes.

FT ECONOMICS Van Bibber et al.33 prepared a process and economic study of a Fischer-Tropsch plant in Healy, Alaska, based on Usebelli coal, a locally mined sub-bituminous coal. Their plant would produce 14,640 bbl/day of FT fluids, of which 58 vol.% would be diesel and 42 vol.% would be naphtha. Diesel prices were assumed to be the same as conventional diesel prices, rather than charge a premium for the higher cetane FT diesel. They also assumed that they could sell the FT naphtha for a $0.10 discount relative to gasoline, rather than dealing with naphtha upgrading issues. We34 modified this report for a Wyoming Powder River Basin (PRB) location. This is appropriate because PRB coal is very similar to Usebelli coal. To account for the higher elevation of the PRB site, another compressor was added to the air separation unit. Capital costs were updated to 2008. Locally, there is a demand for carbon dioxide for enhanced oil recovery, so we assumed that we could sell carbon dioxide for $1.00 per million standard cubic feet. This assumption had little impact on plant economics. With carbon capture and sequestration, the plant would cost about $1.9 billion. The FT fluid cost would have to be $89/bbl for a 10% internal rate of return. This is close to the pre-tax price of gasoline and diesel when this was written. Figure 13.9 is a tornado

Figure 13.9 Effect of a 25% change in assumed economic parameters on the percent internal rate of return for a Fischer-Tropsch coal-to-liquids plant with carbon capture and sequestration in the Powder River Basin of Wyoming. For the base case, $89/bbl is the assumed product price.

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plot showing how a 25% change in assumed economic parameters affects the internal rate of return. Not surprisingly, the plant economics are sensitive to product price. The process economics are dominated by capital construction costs, so changes in capital costs and plant availability have a strong impact on profitability. Coal costs, on the other hand, have a minor effect. The average mine-mouth price of Wyoming coal in 2008 was $11.39/ton, so a 25% change in this price does not have a large impact on operating costs. The price of carbon dioxide has little impact on the overall economics.

REFERENCES 1. Nahas NC. Exxon catalytic coal gasification process, fundamentals to flowsheets. Fuel. 1983;62: 239-241. 2. Great Point Energy. Our technology. ; 2008. 3. Panek JK, Grasser J. Practical experience gained during the first twenty years of operations of the great plains gasification plant and implications for future projects. U.S. Dept. of Fossil Energy, Office of Fossil Energy, ; 2006. 4. Haldor Topsoe. From solid fuels to substitute natural gas using TREMPÔ. . 5. Fischer-Tropsch Archive. www.fischer-tropsch.org. 6. Sasol. 50 Years of innovation. . 7. Anderson RB. The Fischer-Tropsch Synthesis. Academic Press; 1984. 8. Lamprecht D. Fischer-Tropsch fuel for use by the U.S. military as battlefield-use fuel of the future. Energy & Fuels. 2007;21:1448-1453. 9. American Society for Testing and Materials. Standard specification for diesel fuel oils. ASTM D 975e05. 10. U.S. Energy Information Administration. tonto.eia.doe.gov/dnav/pet/pet_pri_refoth_dcu_nus_a.htm. 11. Lu X. Computer aided process design for Fischer-Tropsch conversion of coal to liquid fuels. M.S. thesis, University of Wyoming; 2003. 12. Duvenhage DJ, Shingles T. Synthol reactor technology development. Catal. Today. 2002;71:301-305. 13. Sie ST, Krishna R. Fundamentals and selection of advanced Fischer-Tropsch reactors. Appl. Catal. A: Gen. 1999;186:55-70. 14. Jager B. Fischer-Tropsch reactors, Presented at the Am. Inst. of Chemical Engineers Nat. Meeting, New Orleans; March 31-April 4, 2003. 15. Bartholomew CH. Recent developments in Fischer-Tropsch catalysis. Catal. Lett. 1990;7:303-316. 16. Kilanowski DR. Microchannel Fischer-Tropsch approaches market entry phase. Presented at the 4th BTLtec, Graz, Austria, Sept. 24e25, 2009. . 17. Bell DA, Costopolous SE. A Fischer-Tropsch fluids refinery. Presented at the Am. Inst. of Chemical Engineers Nat. Meeting, Salt Lake City, Utah; Nov. 2007. 18. Harkins SEC. Fischer-Tropsch fluids refinery. M.S. thesis, University of Wyoming; 2010. 19. Sie ST. Acid-catalyzed cracking of paraffinic hydrocarbons. 3. evidence for the protonated cyclopropane mechanism from hydrocracking/hydroisomerization experiments. Ind. Eng. Chem. Res. 1993;32:402-408. 20. Leckel D. Hydrocracking of iron-catalyzed Fischer-Tropsch waxes. Energy & Fuels. 2005;19:17951803. 21. Leckel D, Liwanga-Ehumbu M. Diesel-selective hydrocracking of an iron-based Fischer-Tropsch wax fraction (C15-C45) using a MoO3-modified noble metal catalyst. Energy & Fuels. 2006;20:2330-2336. 22. Leckel D. Selectivity effect of oxygenates in hydrocracking of Fischer-Tropsch waxes. Energy & Fuels. 2007;21:662-667.

Substitute Natural Gas and Fischer-Tropsch Synthesis

23. Leckel D. Low-pressure hydrocracking of coal-derived Fischer-Tropsch waxes to diesel. Energy & Fuels. 2007;21:1425-1431. 24. Krishna R, Joshi GC, Purohit RC, Agarwal KM, Verma PS, Bhattacharjee S. Correlation of pour point of gas oil and vacuum gas oil fractions with compositional parameters. Energy & Fuels. 1989;3: 15-20. 25. Brent FD, Shah L, Berry E, Schrader CH, Anderson J, Erwin J, et al. Early entrance coproduction plant, task phase II, task 2.5: Fischer-Tropsch product upgrading. Texaco Energy Systems, U.S. Dept. of Energy DE-FC26e99FT40658; 2004. 26. Symoniak MF. Upgrade naphtha to fuels and feedstocks. Hydrocarbon Proc. 1980;110-114. 27. Riddler DE, Twigg MV. Steam reforming. In: Twigg MV, ed. Catalyst Handbook. 2nd ed. Wolfe Publishing; 1989;225-282. 28. Krishna KR, Rainis A, Marcantonio PJ, Mayer JF, Biscardi JA, Zones SI. Next generation Isodewaxing and hydrofinishing technology for production of high quality base oils. Presented at the NPRA Lubricants and Waxes Meeting, Houston; Nov. 14-15, 2002. 29. American Petroleum Institute. Appendix E: API base oil interchangeability guidelines for passenger car motor oils and diesel engine oils. ; 2009. 30. Claude MC, Martens JA. Monomethyl-branching of long n-alkanes in the range from decane to tetracosane on Pt/H-ZSM-22 bifunctional catalyst. J. Catal. 2000;190:39-48. 31. Claude MC, Vanbusele G, Martens JA. Dimethyl branching of long n-alkanes in the range from decane to tetracosane on Pt/H-ZSM-22 bifunctional catalyst. J. Catal. 2001;203:213-231. 32. Green CL. U.S. base oil price report. Lube Report. 2010;10. no. 10. . 33. Van Bibber L, Thomas C, Chaney R. Alaska coal gasification feasibility studies-Healy coal-to-liquids plant. DOE/NETL 2007/1251. ; 2007. 34. Schaffers WC, Bell DA. Economics of Wyoming coal gasification. Presented at the Int. Pittsburgh Coal Conf., Pittsburgh; Sept. 21-24, 2009.

391

INDEX Ab initio, 163-4, 239 ACF, see activated carbon fiber Acid gas removal, 116-130, 139, 142-3 Activated carbon, 157, 159, 162-3, 168-9, 170, 177, 183, 187, 196, 203, 206, 237-9, 258, 260, 261, 262, 263, 264, 268, 270, 273, 282-288, 311-314 AX-21, 162-3, 177, 183 bromine impregnated, 262 injection, 257, 258 treated, 259 Activated carbon fiber, 171, 183, 238-9 Air separation unit, 24, 150-1, 154-5, 342, 344 Air pollution control devices, 248, 251, 253, 288 Air toxics, 18 Alanates, 157, 218, 220, 223, 243-4 Alter NRG plasma gasifier, 96-8 Alumina, 142, 144, 265, 344 Amalgamation, 248, 258, 265, 266, 278-281 American Society for Testing and Materials, 2-5, 7 Amines, 22-3, 27, 124-6, 145, 265 adsorbents, 314-318 acid-base reactions carbamate formation, 125 hindered, 125 monoethanol amine, 22-3, 27, 124, 145, 314 primary, 125 secondary, 125 tertiary, 125 Ammonia, xi, 19, 118, 127-8, 145-7, 151, 341-51, 356-7, 378 aqueous, 127-8 as a fuel, 348 catalysts, 341, 344-7 poisons, 345 equilibria, 342-3 history of synthetic, 341, 350-1 hydrogen feedstock, 139, 145-6, 342 synthesis, 341-7 toxicity, 348 Ammonium nitrate, 341, 347, 349-50 Ammonium phosphates, 341, 347 Ammonium sulfate, 127-8, 347-8 Anthracite, 1, 4, 249

APCD, see Air pollution control devices Ash, ix, 3, 5 fluid temperature, 5 fusion temperature, 8 hemispherical temperature, 5, 7 initial deformation temperature, 5, 6 inorganic additives, 8 pozzolanic, 9 softening temperature, 5, 6 temperature of critical viscosity, 7, 8 ASTM, see American Society for Testing and Materials Australian bituminous coals, 56-8 Benfield process, 27, 126-7 Bergius process, 30, 379 BET, see Brunanuer-Emmet-Teller BGL gasifier, 40-1, 75-7 Binding energy, 159, 164, 180, 233-4 Bintulu plant, 379, 384 Biomass, 10 Bituminous coal, 1, 4, 8, 13, 248, 253, 257-259, 263, 271-273 Borohydrides, 157, 218, 225-6, 244 Brunanuer-Emmet-Teller, 169, 170, 173, 182-6, 192, 205 CA, see carbon aerogel Caking of coals, 4 Calcium hydroxide, 61 Calcium sulfide, 130 Carbide-derived carbon, 185, 239 Carbon aerogel, 184, 238 Carbon capture and sequestration, x, 14, 21-7, 134-5, 152-5, 295 co-sequestration with hydrogen sulfide, 134-5 enhanced oil recovery, 134-5 with integrated gasification combined cycle, 152-5 Carbon dioxide, x, 19, 21-28, 38-40, 56-8, 105, 114, 116-130, 137-49, 184, 192, 272, 293339, 353, 358-9, 362-3, 374, 378-9, 389 capture, 301 capture capacities, 302

393

394

Index

Carbon dioxide, x (Continued) emissions, 294 separation processes, 296 sorption, 293-339 Carbon monoxide, 18, 19, 38, 56-8, 105, 137-49, 353-8, 362-3, 373-4, 379 Carbon nanofiber, 157, 159, 168-9, 181-3, 237-8 Carbon nanostructured material, 157, 159, 160, 166, 187 Carbon nanotubes, 157, 159, 160, 162-8, 172-3, 175, 180-1, 187, 235-9 Carbonyl sulfide, x, 114-5, 142 Catacarb process, 126 Catalysts, 53, 59-62, 115, 139-43, 145, 168, 182, 215, 217-8, 249, 252, 267-273, 278, 285, 341, 344-7 ammonia synthesis, 341, 344-7 carbon based, 269 carbonyl sulfide hydrolysis, 115 dimethyl ether synthesis, 361, 364-5 Fischer-Tropsch, 379-80 gasification, 53, 59-62 metal oxide, 269, 271 methanation, 145 methanol synthesis, 353-4, 357-9 nitric acid synthesis, 349 SCR, 269 water gas shift, 139-43 CCS, see carbon capture and sequestration CDC, see carbide-derived carbon CFD, see computational fluid dynamic models Char, 36-9, 56-8 formation, 36-8 oxidation, 38-9 Chemisorption, 157-8, 164, 176, 183, 185, 187, 210-1, 300 Chemistry, gasification, ix, 36-69 Chevron Isodewaxing process, 388 Chevron WWT process, 116 Chromium oxides, 140-1 Clathrate, 157, 207, 209, 210, 242 Claus process, 116, 130-2, 134 Claussius-clapeyron equation, 179 Climate change, 294 CNF, see carbon nanofiber CNTs, see carbon nanotubes CO, see carbon monoxide CO2, see carbon dioxide

Coal, 237, 247-250, 253, 257-259, 261-273, 277-279, 287-288 analysis, 2-4 blending, 248, 249, 257, 288 cleaning, 253, 257, 288 geologic origin, 1 macerals, 2 porosity, 2, 3, 9, 10 price, 13-4 production, 11-3 rank, 4, 5, 52-3, 247, 258, 288 effect on required gasification conditons, 52-3 reserves, 11-14 resources, 11-14 Coalification, 1, 4 Coal-to-liquids, 28-33 Cobalt, 142, 379 Complex hydride, 157-9, 218, 222, 224, 228-9, 243-4 Computational fluid dynamic models, 67-8 ConocoPhillips E-Gas gasifier, 88-9 Copper, 142, 201 Copper oxide, 142 COS, see carbonyl sulfide Crude oil, see petroleum Cryogenic, 158, 167, 169, 203-4, 234 Dehydrated Prussian blue analogues, 195-6, 200, 240 Density functional theory, 163-4, 166, 197, 236 Department of energy, 158, 163, 166, 188, 202, 204, 216, 236 Devolatization, 36-38 DFT, see density functional theory Diesel, 147-8, 370, 380-1, 385-9 Dimethyl ether, xi, 353, 357 catalyst, 361, 364-5 diesel substitute, 360 equilibria, 360 impurity in methanol, 353, 357 LPG substitute, 360 synthesis, 359-85 Direct hydrogenation of coal, 30-33, 149 Bergius process, 30 Catalysts, 30-3 Donor solvent mechanism, 31 Ebullating bed hydrogenation, 31, 149 Exxon Donor Solvent process, 33 Shenhua plant, 30

Index

DKRW Medicine Bow plant, 367-8 DME, see dimethyl ether DOE, see department of energy Double-walled carbon nanotubes, 163 Dry solids pump, 87-8, 95-6 Pratt and Whitney, 95-6 Stamet, 87-8 Drop tube reactor, 63 Durene, 366 DWNT, see double-walled carbon nanotubes

Foster Wheeler Partial Gasifier Module, 52, 79-80 Frasch process, 132 Fuel cell, xi, 146-8, 157, 220, 348 FutureGen, 153

Electrostatic precipitator, 251, 253-256, 278, 288 Encoal, 28-30 Entrained flow gasifiers, 5, 38, 52-3, 65-6, 83-96 ESP, see Electrostatic precipitator Excess adsorption, 157, 159, 160-1, 184 Explosives, 341, 350 Exxon Catalytic Coal Gasification process, 61 ExxonMobil Flexsorb process, 125-6 Exxon Mobil MTG process, 365-8

Gasoline, xi, 147-8, 365-7, 380-1, 386-7 Gas to liquids, 379 GCMC, see grand canonical monte carlo GE gasifier, 83-8, 95 Global warming, xi, 4, 14 GNF, see graphitic nanofiber Grand canonical monte carlo (GCMC), 160, 162-3, 197, 202 Graphitic nanofiber, 163, 166-7, 181-3 Great Plains Synfuels plant, 121-2, 134, 375, 377 Great Point Energy Bluegas process, 62 Greenhouse gas, xi, 4, 18, 153, 294, 378 Gross calorific value, 2 GTL, see gas to liquids Gypsum, 20, 132-3

Feedstocks; air versus oxygen, 41 gasification, ix balancing gasifier feeds, 39 Fertilizer, xi, 132, 341, 347-51 FGD, see flue gas desulfurization Fischer-Tropsch, 110, 155, 379-90 Anderson-Schultz-Flory distribution, 380-1 catalysts, 379-80 economics, 389-90 liquids, 110, 385-9 reactors, 382 advanced Synthol, 382 ARGE, 383-5 microchannel, 384-5 SMDS, 384 Synthol, 382-3 synthesis, 155, 379-85 Fixed carbon, 3 Flat flame burner, 64 Flue gas desulfurization, 20, 133, 253, 258, 264 Fluidization regimes, 64-7 bubbling regime, 65 fast fluidization, 65 fixed bed regime, 65 incipient fluidization, 65 pneumatic conveying, 65 Fluidized bed gasifiers, 5, 38, 52, 77-83, 153

H2, see hydrogen Haber-Bosch process, 341 Haldor Topsoe TIGAS process, 365 Haldor Topsoe TREMP process, 376 Heat of adsorption, 166, 186, 196-7, 200, 202-3, 233-5, 241 Heat transfer materials, 40 Heuristic design of gasifiers HHV, see higher heating value Higher heating value, 2, 3 High pressure, 157-8, 167-8, 187, 191, 201, 204, 209, 232, 238, 243 High-resolution transmission electron microscopy, 176-7 High Temperature Winkler gasifier, 79 Home heating and cooking, 17 HRTEM, see high-resolution transmission electron microscopy H2S, see hydrogen sulfide Hybrid materials, 318 Hydrocarbons, xi, 182 Hydrogen, x, 38-9, 105, 137-49, 157-245, 341-4, 347-8, 353-4, 356, 358-9, 362-3, 369, 373-4, 376, 379 Hydrogen adsorption, 157, 159, 161-173, 175-7, 179-187, 189, 190-5, 197-207, 209, 211, 213, 215, 217-9, 221, 223, 225, 227, 229, 231, 233, 235-243, 245

395

396

Index

KBR Transport gasifier, 44-5, 80-3 Kinetics, 158-9, 210, 216-222, 226, 229, 230, 243-4 Kubas-type interaction, 230, 233, 235

MEA, see amines, monoethanol amine Membrane wall in gasifiers, 91, 93 Mercury, x, 247-289 Oxidation, 250, 266-273, 275, 277 Removal, 251-286 Sorption, 248, 249, 260, 264, 265, 268, 271, 273, 279, 281, 282, 284, 287 Species, 249, 278 Mesoporous carbon, 184-6, 239 Mesoporous titanium oxide, 234, 245 Metal hydride, 157-9, 211, 213-4, 216-9, 230, 242-4 Metal organic framework, 157, 186, 188-204, 207, 230, 235, 239-241 Metaplast, 36-7 Methane, 38-9, 43-5, 61-2, 105, 145, 239, 373-9 equlibria, 374-5 steam reforming, 43-5 substitute natural gas, 373-9 Methanol, xi, 117-22, 155, 353-70, 378-9 equilibria, 354-6, 361-4 dimethyl ether feedstock, 359-370 olefin feedstock, 368-70 Rectisol solvent, 117-122 synthesis, xi, 155, 353-9 transportation fuel, 359 Micropore volume, 175, 184, 186-7, 194, 199, 202, 205 Microporous titanium oxide, 231, 233-4, 245 Mineral matter in coal, 383 Mitsubishi Heavy Industries gasifier, 92-3 MOF, see metal organic framework Moisture in coal, 2, 3, 273 Molybdenum, 142 Motunui plant, 365-7 Moving bed gasifiers, 5, 37-8, 65, 73-7, 153 Multi-walled carbon nanotubes, 164, 168, 174-8, 180-1, 183, 187, 238-9 MWNTS, see multi-walled carbon nanotubes

LHV, see lower heating value Li, see lithium Lignite, 1, 4, 8, 253, 258, 270, 272, 273, 278 Limestone, 20, 114, 130 Lithium, 168, 207, 218, 226, 229, 231-4, 244 Lithium zirconate, 322 Lower heating value, 2 Lubricating base oil, 387 Lurgi gasifier, 5, 73-5, 121-2

N2, see nitrogen Naphtha, 122, 148-9, 380-1, 386-7, 389 Natural gas, ix, 14, 134-5, 149-52, 242, 373-9, 377-8 combined cycle plant, 149-152 competition with coal gasification, 377-8 liquids, 387 liquified, 377-8 shale, 378

Hydrogen adsorption capacity, 163, 166, 184-6, 191, 199, 204-5, 207 Hydrogen adsorption isotherm, 162, 177, 179, 183, 192, 195, 202 Hydrogen binding enthalpy, 197 Hydrogen chloride, 279 Hydrogen peroxide, 277 Hydroge storage, 157-160, 162-170, 173-8, 180-5, 187-9, 191, 193, 196-207, 209, 211, 214-6, 218-9, 221, 223-5, 229-245 Hydrogen storage capacity, 162-9, 170, 175-8, 180-3, 187-8, 191, 193, 197, 199, 200, 205, 207, 209, 216, 218-9, 221, 223-4, 229, 232, 234, 236-9 Hydrogen storage isotherm, 231 Hydrogen sulfide, x, 22, 114, 116-130, 139, 279 Hydrotalcite-like compounds, 331 Hypercoals, 53 IGCC, see integrated gasification combined cycle Illinois no. 6 coal, 40, 152 In-situ coal gasification, see underground coal gasification Integrated gasification combined cycle, xi, 14, 149-55 International Standards Organization, 8 Iron, 143-5, 271, 277, 344-5, 379 Iron oxides, 140-1, 143-5, 344 Iron sulfide, 141, 145 ISO, see International Standards Organization Isomerization, 385-8 Isosteric heats of adsorption, 179-180, 186, 194 Jet fuel, 370, 387

Index

sour, 134-5 steam reforming, ix substitute natural gas, 373-9 Nickel, 145, 200-1, 241, 379 Nitric acid, 349-50 Nitrogen, xi, 151, 171-172, 184, 192-3, 199, 341-5, 347-9, 351 NOx, 18-20, 26, 151, 250, 252, 263, 266, 269, 270, 273 Nuclear power, ix, 14, 40 Olefins, xi, 368-70, 385-6 Oil, see petroleum Oil sands, 134 Organic foods, 351 Oryx GTL plant, 379 Osmium, 341 Oxidation, 10, 163, 230, 233 Oxycombustion, 24, 154, 258, 264, 277, 287 Oxygen, 10, 24, 38-40, 189, 197, 199, 205-6, 208, 216, 282, 285, 286 Particle density, 160 Particulates, 18, 252 Pearl GTL plant, 379 Peat, 1, 4 Petrography, 2 Petroleum, xi, 11, 14, 31-2, 351, 378, 380, 387 farm use, 351 hydrogenation, 31-2 Petroleum coke, 52-3 Phosphoric acid, 132-4, 347 Physisorption, 157-160, 162, 168, 183, 185, 188, 202, 207, 211, 230, 233, 235, 237, 240, 297 Pittsburgh no. 8 coal, 89 Plasma torch, 97 Platinum, 147, 187-8, 203-4, 217, 230, 239, 243, 245, 268, 272, 279, 281, 349 Pollutants, 17-20, 257, 280 Pore size, 160, 170, 184-8, 192, 196, 198-9, 201-2, 204, 230, 235, 241 Pore volume, 159, 160-2, 185, 187-8, 192-3, 197, 199, 201-2, 232, 235, 241 Porosity, 163, 183, 193, 239, 259 Porosity of coal, 2, 3, 9 Porous material, 157, 159, 186, 202, 239, 240-1 Potassium bicarbonate, 27, 126 Potassium carbonate, 27, 60, 126

Potassium oxide, 344 Powder River Basin coal, 3, 13, 29, 44, 89, 113, 249, 257, 263, 389-390 Powerspan process, 127-8, 347-8 Pratt and Whitney Rocketdyne gasifier, 93-96 Pressure, estimation from sygnas composition, 41-5 Pressure swing adsorption, 297 Pressurized fluidized bed combustion, 27 Process goals for gasification, 35 Proximate analysis, 3, 4 Pulverized coal combustion, 14, 19-24, 152, 253, 263 power cost, 152 supercritical, 20-1 Pyrites, 113, 248, 257 Pyrolysis, 28-30, 36-38 Raman spectra, 164 Reaction rates, 46-64 ammonia synthesis, 345-7 effectiveness factor, 345-6 Thiele analysis, 346 gasification, 46-64 activation energies, 59 catalysts, 53, 59-62, 130 char morphology, 50-1 changing solid mass, 53-6 effectiveness factor, 49-51 experimental techniques, 62-4 Langmuir-Hinshekwood equations pore size effect on mass transfer rates, 48-9 random pore model rate-limiting steps, 46-8 shrinking core model Thiele analysis, 49-51 transition from surface limited to mass transfer limited, 51-3 methanol synthesis, 357-9 mercury oxidation, 275 water gas HT shift, 141 Rectisol process, 117-122, 375 Refractory, 8, 76, 80, 84 Ruthenium, 345, 379 Sargas, 27, 126 Sasol, 379 Sasolburg plant, 379 Scanning electron microscopy, 221

397

398

Index

SCR, see selective catalytic reactor Secunda plant, 379 Selective catalytic reactor, 19, 252, 253, 269, 271, 288 Selexol process, 122-124, 153 SEM, see scanning electron microscopy Shell Claus Offgas Treatment process, 131-2, 134 Shell gasifier, 40-1, 90-91, 95 Siemens gasifier, 92-3 Silicon tetrafluoride, 133 Single-walled carbon nanotubes, 162-7, 169, 172-4, 177-180, 235-8 Slag, ix, 5, 39, 91-2 Slit pore, 160, 162-3, 235 Slurry bubble column reactor, 365 Slurry feed to gasifiers, 86 SNG, see substitute natural gas SO2, see sulfur dioxide Sodium carbonate, 61, 126 Sodium nitrate, 341 Solar heat, 40 Solar power, 152 Soot, 37 SOx, 17, 284 South African coals, 8 Spalling, 106 Spillover, 157, 187, 202-3, 207, 239, 241 Spontaneous combustion of coal, 10 Stamet pump, 87-8 Steam, see water Sub-bituminous coal, 1, 4, 8, 13, 19 Substitute natural gas, 155, 373-9 Sulfur, ix, 3, 113-35, 248, 252, 257 market, 132 Sulfur dioxide, 17, 20, 25-6, 113-4, 130-2, 250, 252, 261, 268-270, 273, 276, 282-287 Sulfuric acid, 132-3 Sulfur trioxide, 132 Surface area, 159, 162-3, 168-9, 170-3, 177, 180-193, 195, 197, 199, 200-2, 204-5, 217, 230, 232, 235, 237-9, 241 SWNT, see single-walled carbon nanotubes Syngas, ix, 278, 279 estimating composition based on equilibrium, 41-5 Tar, 36-38, 75-6, 86, 104 TEM, see transmission electron microscopy

Temperature, estimation from syngas composition, 41-5 Temperature programmed desorption, 166, 177, 201, 224-5 Texaco gasifier, 83 TGA, see thermo-gravimetric analyzer Thermal Power Research Institute gasifier, 93 Thermo-gravimetric analyzer, 62 Thief carbon, 262-264, 271, 288 Total uptake, 157, 159-160, 231-2 TPD, see temperature programmed desorption Transmission electron microscopy, 176, 178, 181 Transportation fuels, xi U-gas gasifier Ultimate analysis, 3 Underground coal gasification, ix, 14, 101-110 concept, 101 connections between injection and production wells, 102-3 controlled, retractable injection point, 101 directional drilling, 102-3 Gastech study, 108, 9 Hanna tests, 108-9 modeling, 103-6 motivation, 102 process control, 103 recoverable coal, 108 reverse combustion technique, 102 water contamination, 107 UOP/HYDRO MTO process, 368-370 Urea, 118, 341, 34-9 Usebelli coal, 389 Utah sub-bituminous coal, 49 Vacuum swing adsorption, 299 Volatiles, 3 Volatile organic compounds, 18 Volumetric storage density, 159, 192, 204 Wabash River Coal Gasification Repowering Project, 89 Water, 1-3, 38-41, 56-8, 84-6, 92-3, 105, 107, 114-6, 129, 137-4, 157, 183, 192-4, 198-9, 201-2, 207-9, 211, 249, 250, 259, 272, 273, 282, 283, 353-4, 358-9 black, 115-6 coal moisture, 1-3, 273

Index

formation in gasifier, 38 hydrogen production, from, x, 42-5, 137-45 methanol synthesis, from, 353-4, 358-9 reaction with carbonly sulfide, 114-5 reaction with char, 38-41, 56-8 reaction with metal sulfides, 129 sour, 115-6 syngas quench, 84-6, 89, 92-3, 115-6 underground coal gasification, 105, 107 Water gas shift, x, 42-5, 137-43, 153, 355, 359, 362 catalysts, 139-43 iron-base HT shift, 139-41 LT shift, 139, 142, 153 sour, 142-3 equlibria, 42-5, 138 methanol synthesis catalyst activity, 355, 359, 362 sour gas shift, 142-3, 153

Wax, 381, 385-8 FT, 381, 385-6, 388 paraffin, 381, 387-8 Weathering of coal, 10 West Virginia coal, 23 Weyburn enhanced oil project Wind power, 14, 152 Winkler gasifier Wire mesh reactor, 64 ZeroGen, 153 Zeolite, 157, 168, 186, 190, 204-7, 230, 239-242, 303-311 Zeolite-like carbon, 186, 239 Zinc chloride, 259, 260 Zinc ferrite, 129 Zinc oxide, 128, 142 Zinc sulfide, 129 Zinc titanate, 129

399