R - 3 r 7 LA-UR-97- 7 1 2

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CON F- 9 2 0 5 aa - - &- A METHOD FOR PERMANENT DISPOSAL OF CO, IN SOLID FORM

DARRYL P. BUTT, MST-6 ALANA BENJAMIN, MST-6 TERRY G. HOLESINGER, MSTB YOUNGSOO S. PARK, MST-6 KLAUS S. LACKNER, T-3 CHRISTOPHER H. WENDT, T-3 ROBERT P. CURRIER, ESA-EPE DAVID M. HARRADINE. CST-6 MEREDITH RISING, MlT KOJl NOMURA, CHlCHlBU ONADA CEMENT COMPANY

GLOBAL WARMING INTERNATIONAL CONFERENCE AT COLUMBIA UNIVERSITY, NY IN MAY 1997

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Form No. 836 R5 ST2629 1w91

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F

A Method for Permanent Disposal of CO, in Solid Form

Darryl P. Butt, Klaus S. Lackner, Christopher H. Wendt, Alana Benjamin, Robert Currier, David M. Harradine, Terry G. Holesinger, Youngsoo S. Park, and Meridith Rising

P. 0. Box 1663, M. S. G755 Los Alamos National Laboratory, Los Alamos, NM 87545

505-667-9307

Koji Nomura Chichibu Onoda Cement Co., Tokyo, Japan

Addresses of key authors: Darryl P. Butt P. 0. Box 1663, M. S. G755 Los Alamos National Laboratory Los Alamos, NM 87545 505-667-9307

Klaus Lackner P. 0. Box 1663, M. S. B216 Los Alamos National Laboratory Los Alamos, NM 87545 505-667-5694

Christopher Wendt 31 Bagel Ct. Madison, WI 53705 608-23 1- 13 13

A Method for Permanent Disposal of Carbon Dioxide in Solid Form

Key Words: carbon dioxide, green house effect, serpentine, magnesium hydroxide, carbonation

Summary: We describe a method for binding the greenhouse gas carbon dioxide as magnesium

carbonate, a thermodynamically stable solid, for safe and permanent disposal, and with

minimal environmental impact. The technique is based on extracting magnesium

hydroxide from common ultramafic rock for thermal carbonation and subsequent

disposition. The economics of the method appear to be promising, however, many details

of the proposed process have yet to be optimized. Initial estimates indicate that binding

and disposal would impose a burden o approximately 3$/kWH onto the cost of

electricity.’ Realization of a cost effective method requires development of optimal

technologies for efficient extraction and thermal carbonation. In this paper, we describe

-6 cptd ?w t r k d ~ +e f uat[\l r- +L c b - j 4 . m &bra-b n r r ~ m#bh 9

V&t.P<

some of the kinetic limitations and opportunities. The proposed disposal technique may

be viewed as a sort of insurance policy in case global warming, or the perception of

global warming causes severe restrictions on CO, emissions.

2

Introduction

There is a great deal of evidence that CO, levels in the atmosphere affect the

environment and are increasing at a potentially dangerous rate. Today, the level of

carbon dioxide in the atmosphere is 30% higher than at the beginning of the 19th century,

with half of this increase occurring in the last 30 years (Siengenthaler, 1987; Keeling,

1995). Consequently, there is little doubt that this increase is due to anthropogenic

causes. However, the potential consequence of this rise is and will continue to be argued

for years to come. There is a significant number of researchers who predict highly

deleterious effects of carbon dioxide, in particular, on the global climate (Manabe, 1967;

Ramanathan, 1988; Darmstadter, 1989; Schneider, 1989; Hidy, 1994).

Regardless of whether or not the more ominous predictions will turn out to be

correct, they are sufficiently alarming to motivate the search for alternative means for CO,

disposal. If society could develop ways to safely and economically dispose of CO, in

thermodynamically stable forms, we could eliminate the potential danger of climatic

changes due to greenhouse gas emissions from burning of fossil fuels. Toward this end,

we have recently outlined and developed a methodology for binding carbon dioxide in a

solid form that is thermodynamically stable at the earth’s surface (Lackner, 1995;

Lackner, 1996; Butt, 1996). The carbonation of Mg(OH), is of particular interest

because of the abundance of magnesium on earth and the relative ease with which it can

be extracted from Mg-bearing minerals. As will be discussed in greater detail below, the

major focus of our research has been on developing ways to extract Mg(OH), from

common ultramafic rock, such as serpentinite and peridotite, and then rapidly and

efficiently carbonating the powder to MgCO,. The quantities of accessible deposits of

3

these minerals vastly exceed the quantities of available fossil fuel in the world. Because

this active approach would result in a waste that is thermodynamically stable, it would be

possible to permanently dispose of great quantities of CO, with minimal environmental

impact and without the danger of a sudden accidental release of gaseous CO,, which, as

will be discussed in more detail below, has proven fatal in even in comparatively small

releases of gas.

The 10,000 Gigatons of coal reserves, in particular, could provide the world with

energy for tens of generations (United Nations, 1993; Schomber, 1993). By

comparison, we currently consume about 6 Gigatons of coal annually (United Nations,

1993; Hidy, 1994). However, coal’s long term use may be severely curtailed if we

continue to dispose of CO, into the atmosphere. The availability of a CO, fixation

technology would serve as insurance in case global warming, or the perception of global

warming causes severe restrictions on CO, emissions. If the increased energy demands

of a growing world population are to be satisfied from coal, the implementation of such a

technology would quite likely be unavoidable. There are now a number of proposed

schemes other than our own for collecting and/or disposing of CO,. For example, ocean

disposal as liquid CO, or clathrites, the use of bacteria in the ocean to collect CO,,

planting of trees, and underground disposal. As with our technology, there is often a

cost penalty for any technology which can reasonably affect global emissions. We

emphasize that there are other techniques that can provide small reductions in CO,

emissions at little cost (Audus, 1995). Thus, we view our CO, fixation concept as a

viable option in case greater measures are required due to political or environmental

pressures.

The economics of binding CO, in a solid form is strongly dependent on the

kinetics of reactions between CO, and the host substrate. The kinetics of simultaneous

4

thermal dehydroxylation and carbonation of precipitated Mg(OH), has been studied in

some detail (Butt, 1996). We demonstrated that during carbonation, MgCO, precipitates

on the surface of disrupted Mg(OH), crystals acting as a kinetic barrier to both the

outward diffusion of H,O and the inward diffusion of CO,. In this paper we describe our

experimental work on extracting and carbonating Mg(OH),. The carbonation kinetics

appear to be rapid enough for an economical, industrial process using fine particle size

distributions and high pressures.

The Importance of the Safe and Permanent Disposal of Carbon Dioxide

In many respects, the disposal or storage of greenhouse gases such as CO,, can

be viewed as having safety issues analogous to those associated with storage of other

hazardous materials. Nature has given us a tragic, but valuable example of the potential

danger of stored CO,. In 1986, there was a massive release of CO, from Lake Nyos,

which sits above a small village in Cameroon, Africa (Kling, 1987; Ladbury, 1996;

Freeth, 1994; Zhang, 1996). Being heavier than air, the cloud of gas flowed down the

mountainside over the village and for distances of approximately 10 km into the

surrounding countryside, asphyxiating all but plant life, including more than 1700 people

and 3000 cattle.

All evidence suggests that the release of CO, was caused by a hydrodynamic

instability. The most plausible explanation for the event seems to be that described by

Bang (Zhang, 1996). He proposed (and demonstrated the feasibility) that, due to a

natural perturbation in the lake, the dense, CO, rich water at the bottom rose to a point

where pco 2 ptotal. Under this condition bubbles nucleated and grew and, subsequently,

rose quickly to the lake’s surface. The force associated with the rising gas lead to an

2

5

overturning of the lake, causing increasingly vigorous outgassing, and ultimately a kind

of limnic eruption of CO,.

The 0.1 km3 of CO,, reported to be released in the disaster, was the equivalent of

a week of emissions from a 1 GW, coal-fired powerplant. Thus, a large accidental

release of CO, from an unstable deposit could be catastrophic. Further, consider the

safety issues that would be associated with storing the output of hundreds or thousands

of powerplants over many decades. Additionally, even slow leaks of CO, could be

problematic since it only postpones, and could, therefore, compound the problem for

future generations. Thus, it has been our objective in this research to offer a very safe

method for binding CO, in a form that is stable virtually forever.

Summary of the Proposed Magnesium Hydroxide Extraction Process

A significant fraction of the world is covered with carbonated rock, mostly rich in

Ca and to a lessor extent Mg. Essentially, what we are proposing is to make use of the

huge quantity of rock in the world that has not been carbonated by natural events.

Magnesium bearing minerals constitute the most abundant and technologically available

source of such material. Un-carbonated, Ca bearing materials are comparatively less

common and we believe would be more costly to process. In particular, peridotites and

serpentinite, which are approximately 30-50% MgO by weight, are available in quantities

far exceeding what is required to consume the world’s anthropogenic CO,. Most of our

research to date has focused on developing processes relying on serpentinite, which in its

purest form has the composition 3Mg0*2Si0,*2H2O, and, therefore, contains 43.6 wt%

MgO.

6

The direct carbonation of serpentinite, as well as other minerals, is relatively

slow. In order to economically bind the CO, in the host substrate, the magnesium must

be extracted from the mineral. Fortunately, methods for extracting magnesium from

minerals were worked out in some detail around the time of World War 11 and the advent

of sea water extraction processes (Houston, 1945; Gee, 1946). The economics of the

processes were never optimized but appeared to be quite viable. Making use of this early

and seemingly forgotten research, we have set out to develop an economical process

whereby the magnesium is extracted by first dissolving the mineral in hydrochloric acid.

Through a multi-step extraction process, Mg(OH), is precipitated out of solution and then

thermally carbonated. Most of the HC1 is recovered and reused in the extraction process.

A simple schematic of the basic material flow is shown in Fig. 1 for a system using a

peridotite with 45.8 wt% MgO. Figure 2 shows a more detail including impurity removal

steps for a more realistic system using a relatively impure serpentine. Technologies for

collecting and transporting CO, are very well understood (Audus, 1995). Thus, the

focus of our work has been on the bottom half of Fig. 1 including the direct carbonation

step. Many of the details of each step and the economics have been discussed

preliminarily elsewhere (Lackner, 1995; Lackner, 1996), and are being described in

greater detail in a forthcoming paper (Lackner, 1997a; Lackner, 1997b).

Table 1, summarizes the thermodynamics and mass balance equations relevant to

the overall process outlined in Figs. 1 and 2. The overall process is exothermic and is

defined by the following net reaction:

Mg,Si,O,(OH), + 3CO,(g) + 3MgC0, + 2Si0, + 2H,O

As written, equation 1 occurs very slowly except at very high pressures and

temperatures. For example, in relatively unoptimized experiments, we have gotten the

7

V L

reaction to go to near 25% completion by exposing 100 pm serpentine powder to 340 atm

of CO, at 500°C for two hours. Thus, in order to make the reaction go to completion

under less costly conditions, we extract a more reactive form of magnesium by first

dissolving the serpentine in hydrochloric acid as follows:

Mg,Si,O,(OH), + 6HCl(aq) + 13H,O + 3(MgC1,.6H2O) + 2Si0, (2)

By heating the solution and distilling off water and HC1 according to the series of

reactions defined in table 1 and outlined in Fig. 2, a mixture of Mg(0H)Cl and HCl is

produced. By adding water back to the solution, we then precipitate out Mg(OH), in

MgC1,*6H20. In the laboratory, we can do this in essentially a two step process as

follows:

- 160°C MgC1,*6H20 A MgCl(0H) + HCl + 5H,O (3)

followed by:

(4) H2O 2MgCl(OH) ___j Mg(OH), + MgCl,

The Mg(OH), is easily filtered from the solution for carbonation. In the following

section, we describe the end process; that is, the essential and perhaps critical step of

converting Mg(OH), to a stable carbonate.

8

The Direct Carbonation Step

Preliminary studies of the carbonation kinetics of Mg(OH), have been completed

under atmospheric conditions (Butt, 1996) as well as a limited number of studies at 30 to

440 atmospheres (Butt, 1997). The results generally follow trends predicted from

thermodynamics. However, the processes can be quite complicated and the mechanisms

are not fully understood at this point. Understanding the mechanisms could lead to

significant improvements in the efficiency of the process. Figure 3, which was calculated

from thermodynamic data, shows the effect of pressure on the dissociation temperatures

of MgCO, and Mg(OH),. It is apparent that the dissociation temperatures of both the

carbonate and hydroxide increase significantly with pressure. Thus, the carbonation of

Mg(OH), is inherently limited by the dissociation temperature, which is a function of the

pressure of carbon dioxide. Above the dissociation temperature, which is near 400°C at

atmospheric pressure, the carbonate will decompose according to the reaction:

MgCO, + MgO + CO,(g) ( 5 )

Our studies have shown that there is a strong interconnectivity between the

dehydroxylation and carbonation reactions. As illustrated in Fig. 4, the rate of

carbonation is most rapid very near the dissociation temperature of MgCO, (near 390°C at

the altitude of our laboratory). Note that the data in Fig. 4 were determined by measuring

the CO, evolved during reaction with 6M HC1 (Butt, 1996; Pile, 1997). During

carbonation of Mg(OH),, there is simultaneous dehydroxylation and carbonation. Thus,

it is important to fully understand both the kinetics and mechanisms of each process. The

kinetics and mechanisms of dehydroxylation of single crystal Mg(OH), are relatively well

understood (Gordon, 1966a; Gordon, 1966b). In the case of precipitated Mg(OH),

9

(which is polycrystalline) the kinetics of dehydroxylation can vary significantly from one

material to another (Gregg, 1949; Anderson, 1962; Gordon, 1966a; Gordon, 1966b;

Laureiro, 1991; Halikia, 1993; Butt, 1996). Also, the mechanism of dehydroxylation of

precipitated Mg(OH), has not been fully elucidated. The variations between the different

studies appear to be real; that is, the kinetics are strongly dependent on purity, sample

size, crystallite and agglomerate size, thermomechanical history, and environment.

Despite these complications, we can still make some general statements about the process

of simultaneous dehydroxylation and carbonation.

As shown in Fig. 4, using a 20 pm precipitated Mg(OH), powder, the

carbonation reaction reached approximately 20% completion in 30-60 minutes, assuming

that the carbonate that forms is MgCO,. We have been able to get substantially more

carbonation (approximately 50% completion) under the same conditions by grinding our

powders to a finer particle size. However, the efficiency of carbonation at atmospheric

pressure is relatively slow. This is due to the fact that the carbonate that forms is a barrier

to both the outward diffusion of water, thus inhibiting dehydroxylation, and the inward

diffusion of CO,, thus inhibiting carbonation (Butt, 1996). This is illustrated by the X-

ray diffraction data (XRD) shown in Fig. 5. The sample that was heated to 350°C in

helium was converted almost entirely to MgO, while samples heated in CO, retained a

significant amount of Mg(OH), up to the MgCO, dissociation temperature.

Typically, the rate of a thermally activated reaction increases with temperature.

As indicated in Fig. 3, by increasing the pressure of CO,, we can substantially increase

the temperature at which MgCO, dissociates. Therefore, by increasing the pressure of

CO,, we can increase the temperature at which we perform the carbonation step and

presumably accelerate the carbonation rate. We have done this and found that relatively

modest increases in pressure, give substantially more rapid rates of carbonation. For

10

example, at 565°C and 52 atm, we’ve been able to get the reaction to go to approximately

90% completion in 30 min. using the same 20 pm powder described above. The final

product is significantly different from that obtained at lower pressures and temperatures.

As shown in Fig. 9, the crystallites within the particles are transformed to a mixture of

MgCO, with a minor amount of MgO. Also, the final product is more chemically stable

compared with material heat treated at lower temperatures. This is illustrated in Fig. 10

where we show the rate at which CO, is liberated upon reaction with 6M HC1. The

powders heat treated at 540 to 565°C react relatively slowly with the acid compared with

samples treated at 430 to 455°C.

Based on these results it appears that the final process will require the use of

moderately high pressures. We do not view this as a problem because CO, is typically

transported under similar pressures. Thus, there may be no need to add extra energy to

the system since the pipeline will supply at least part of the required pressure.

Economics and Niche Markets

Binding our anthropogenic CO, may sound like a colossal undertaking and

difficult philosophical change for society; however, it may be possible to bootstrap our

way to the final solution by first taking advantage of niche markets. For example, in our

operation we can accept and process waste asbestos, a hazardous material that contains

approximately 40 wt. % MgO. Thus, it would be possible to charge a fee to dispose of

this waste making the operation more profitable. Or, in the short term, some fraction of

the final, carbonated products, as well as the silica and iron impurities, could have some

value. For example, the raw materials could be sold for use in glass, concrete, or other

building products. We estimate that in the short term, sales of iron could effectively

reduce the cost of CO, collection by approximately $10/ton (Lackner, 1997a; Lackner ,

11

1997b). Other niche markets could reduce the process costs to as little as l# versus our

more conservative estimate of 3#/kWh (Lackner, 1995).

However, in order to consume a sufficient quantity of CO,, we will in all

likelihood need to ignore the niche markets in our final economic analysis because of the

great quantities of material that will necessarily be generated. That is, the value added to

the final, carbonated product will be negligible. Thus, disposal ultimately becomes our

only real option. As noted above, we are in the process of doing a rigorous economic

assessment of our process, using the data that we have available, and will report those

results in at least one forthcoming publication (Lackner, 1997a; Lackner, 1997b).

However, in a previous paper we assessed the probable cost of our process by

comparing it to other, simdar large-scale mining and chemical operations (Lackner,

1995). In these estimates we ignored the cost associated with extracting and shipping the

CO, because these are upfront costs that must also be incurred by other competitive

technologies that aim at removing CO, by an engineered route. Using relatively simple

but logical assumptions, it was shown that the cost of binding CO, by our process would

be on the order of $30 per ton of carbon dioxide. This translates into about 3$ per kwh

based on the U. S. emissions of CO, from coal, which amount to approximately 90 gms

of CO, per MJ of produced energy (EIA, 1994). Thus, if implemented immediately, our

process would roughly double the cost of electricity. However, the efficiency of coal

fired processes, in particular, has steadily improved during the last 30 years. Thus,

through relatively gradual implementation of our technology and continued improvements

in efficiency, the consequence would more likely be that the cost of electricity would

remain nearly constant rather than decreasing with time.

12

Summary Remarks and Future Directions for Research

Regardless of one’s view of the consequences of global warming, it would seem

prudent for the U. S. and other countries to invest heavily in research aimed at

understanding the both the effects of and how to better mitigate CO, emissions. The

Japanese, for example, have recognized the importance of CO, mitigation technologies

and recently established the Research Institute of Innovative Technologies for the Earth

(EUTE) which is an organization largely dedicated to studying ways to collect, bind, and

reuse CO,.

Our research has been aimed at binding CO, in a thermodynamically stable solid,

namely magnesium carbonate. In our laboratory we have demonstrated the feasibility of

extracting Mg(OH), from common minerals; and that the Mg(OH), can be converted by a

direct thermal treatment to a stable carbonate for subsequent disposition using CO,

pressures comparable to those used in its normal transportation. The kinetics of

extraction and carbonation have yet to be optimized. Likewise, thermal management in

the Mg(OH)2 extraction process has yet to be optimized. Further research at both our

laboratories (Los Alamos and Chichibu Onada Cement Co.) is aimed at enhancing the

reaction kinetics. However, we are nearing the stage where a pilot demonstration of the

entire process can be justified. It is hoped that the first such system will be developed

within the next few years.

The processes we have developed can be carried out on a massive scale, sufficient

enough to bind in theory all anthropogenic CO,. The quantity of useful minerals in the

world, namely serpentinite and peridotites, are far in excess of what is needed to collect

the CO, emissions from the remaining world’s supply of coal.

13

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14

(Hidy, 1994) G. M. Hidy and D. F. Spencer, “Climate Alteration,” Energy Policy, 22, 1005-1027 (1994).

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15

t c ,

Geologicd Survey Bulletin, 1452, Washington, D. C. (1978), reprinted with corrections, 1979.

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16

Table 1. Heat balances for reactions important to the extraction of Mg(OH), from serpentine and its subsequent carbonation. Note equation 1 is the net reaction of equations 2 and 12,2 is the net of 3 an 4,4 is the net of 5 and 11,5 is the net of 6 and 7, and 8 is the net of 9 and 10. Thus, the molar heat of reaction 1 equals the sum of 2 and 12, and so on. The data were obtained from several sources (Robie, 1979; Chase, 1985; Barin, 1993).

Eqn. No. Reaction Heat/ Cor . Factor Hedmole CO, Equation Hlmol

1 Mg,Si,O,(OH), + 3C0, -+ 3MgC0, + 2Si0, + 2%0(1) -190.8 113 -63.6 2 Mg,Si,O,(OH), + H,0(1) + 3Mg(OH), + 2Si0, 52.5 113 11.5

4 MgC12*6H,0 -+ Mg(OH), + 2HCl(aq) + 4&0(1) 96.7 111 96.7 5 MgC1,*6Hz0 + MgCl,*H,O + 5%0(1) 103.3 2/l 206.5 6 MgC1,.6H,O + MgCl,*H,O + 5H20(g) 323.3 2/1 646.6

3 Mg,Si,O,(OH), + 6HCl(aq) + 13H,0(1) + 3(MgC1,*6H20) + 2Si0, -234.3 113 -78.1

7 HzO(g) + & W ) -44.0 10/1 -440.1 8 MgCl,*H,O + Mg(0H)Cl + HCl(aq) -0.15 2/1 -0.3 9 MgCl,*H,O + Mg(0H)Cl + HCl(g) 74.1 W l 149.4 10 HCl(g) + HCl(aq) -14.9 W l -149.7 11 2Mg(OH)CI + 6H,O + Mg(OH), + MgCl,*6H,O -109.5 111 -109.5 12 Mg(OH), + CO, + MgCO, + H,0(1) -81.1 111 -81.1

17

Power Plant Earth Moving

Carbon Dioxide Scrubbing and

MgC03 Direct 46 ktonslday

t Peridotite MgC1, -b HClExtraction -I

1 hour residence 52 ktons/daY Peridotite Mine

4- Recovery - -26 ktonslday Acid Return

MgC4 0.5 tondday

(HCl Make Up)

Figure 1. A simple diagram showing the flow of material in a CO, disposal scheme for a process using peridotite as the source for magnesium. The processing rates are geared to eliminate the emissions of a coal fired power plant producing 1 GW of electric power at 33% conversion efficiency. The focus of our research has been primarily the bottom have of the diagram including the direct carbonation step.

1

Figure 2. A diagram showing the flow of material in a CO, disposal scheme using relatively impure serpentine as the source for magnesium. Following dissolution of impure serpentine in concentrated HC1 (mix-1), the insoluble SiO, is removed. Recycled MgClOH is then used to precipitate Fe by raising the pH (mix-2). The resulting liquid stream is then combined with recycled MgCl, (aq) and residual brine from the evaporators. This mixture then passes through a series of evaporators (units E-1 to E-3) which vaporize the HC1 and water. The resulting precipitate (MgCIOH) is removed and the concentrated brine is recycled. MgClOH produced in the evaporator units is subsequently re-dissolved in water (mix-4) to produce the Mg(OH), product and MgCl,(aq) for recycling. The Mg(OH), is then used for direct carbonation, which has been the primary focus of our research. We note that a portion of the water produced in the carbonation units (carb-H,O) is combined with purest water stream from the evaporators to dissolve the MgClOH. This permits discharging excess water from the cleanest stream available, i.e. the water produced in the carbonation units.

2

100% complete in less than 30 mi111A

8

53 CI

.r( Y

.r( cb 0

m .r(

200 0.1 1 10 100 1 000 io4

Pressure, atm

Figure 3. Calculated thermodynamic dissociation temperatures for MgCO, and Mg(OH), to MgO and CO, or H,O, respectively. The arrows show how the efficiency of carbonation improves with pressure and temperature for a precipitated Mg(OH), powder heat treated 30 minutes (Butt, 1996; Butt, 1997). Note the x-axis refers to the total CO, or H,O pressures.

3

0 " " " " " " " ' ~ " ~ " ~ " 0 100 200 300 400 500

Temperature, "C

Figure 4. Measured CO, content of precipitated Mg(OH), carbonated for 30 and 60 minutes in 0.765 atm of CO,. The powder had a 20 pm average particle or agglomerate size, where the agglomerates were made up of approximately 100 nm Mg(OH),.crystallites. Note these data were obtained by measuring the volume of CO, liberated during reaction with 6M HCl as described in detail elsewhere (Butt, 1996; Pile, 1997).

3000 1

2500 r

.d a 2000 1

B

$ 1500 : d U

.d * d id

2 1000 :

- 500

Mg(OH&- MgO =o MgCO,=o As Received - I I I 0

0 20 40

375"C, COJ x x

350°C, CO

60

Angle, 28

80 100

Figure 5. X-ray diffraction patterns of as-received Mg(OH), powder and those after 12 hour heat treatments in 0.765 atm of carbon dioxide and helium.

5

2500

2000

500

0

315"C, 0.765 am

0 20 40 60 80 100 Angle, 20

Figure 6. X-ray diffraction patterns of Mg(OH), powder carbonated in 0.765 to 52 atm of carbon dioxide.

6

Figure 7. Transmission electron micrograph showing the crystalline morphology inside an agglomerate of precipitated Mg(OH), The platelet shaped crystallites are Mg(OH),. There is a slight diffraction ring assoclated with MgCO, which is present as an impurity (=2 wt. %) in the powder.

7

Figure 8. Transmission electron micrograph showing the crystalline morphology inside an agglomerate of precipitated Mg(OH), after carbonation for 12 hours in 0.765 atm of CO, at 375°C. What’s left is a collection of disrupted Mg(OH), crystallites in which MgO has precipitated. Very little MgCO, can be detected by TEM or XRD, although from gas analysis it was determined that the sample contained approximately 6% CO, by weight.

8

.

Figure 9. Transmission electron micrograph showing the crystalline morphology inside an agglomerate of precipitated Mg(OH), after carbonation for 30 minutes in 52 atm of CO, at 565°C. What's left is a collection of nanometer sized MgCO, crystallites with a minor amount of unconverted MgO. From gas analyses it was determined that the sample was 42% CO, by weight, that is, the carbonation reaction was nearly 90% complete.

9

c

Mg(OW2 carbonated 30 min in CO, I " " l ' " ' l " " l " ' ' -

Theoretical COz content of MgC03-3 50 - -

r 540"C, 30 atm, medium flow

3

L5 > 40 - 3 0

8" 30 @ 2 M -- 20 $

10

0

565°C. 52 am, high

455"C, 52 atm, low flow

430°C, 30 atm, medium flow

455OC, 52 atm, high flow

0 10 20 30 40 50 Time of Reaction with 6 Molar HCl, min

Figure 10. Plot of weight percent carbon dioxide evolved during reaction with 6M HCl from Mg(OH), samples carbonated between 30 and 52 atm at various temperatures.

10