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Removal of inorganic compounds via supercritical waterLeusbrock, Ingo

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This chapter has been published as:Leusbrock, I., Metz, S. J., Rexwinkel, G., and Versteeg, G. F.; The solubility of magne-sium chloride and calcium chloride in near-critical and supercritical water ; The Journalof Supercritical Fluids 53(1-3), 17-24.

Chapter 5

The solubility of magnesium chlorideand calcium chloride in near-criticaland supercritical water

- 109 -

Chapter 5 ∥ Magnesium Chloride and Calcium Chloride

Abstract

Applications using supercritical water often encounter the presence of

inorganic compounds in feed streams, most often with a minor concentration.

These compounds can lead to damage of the equipment via erosion, scaling

and corrosion or can influence and disturb the main reaction and processes

inside the systems. In order to avoid these problems and to predict the

influence of these compounds, it is vital to posses knowledge of the properties

of the most common inorganic compounds in supercritical water.

In continuation of earlier works of the authors, the solubilities of MgCl2

and CaCl2 are investigated via a continuous flow method in the range of

660 to 690 K and 18.5 to 23.5 MPa. Contrary to earlier experiments

with single-valent salts, precipitates were found during the experiments with

MgCl2 after cleaning the setup. These precipitates were analysed via EDX

and ATF-IR. In the course of the experiments, a decrease in pH of the

samples was investigated what was caused by a parallel hydrolysis reaction.

The solubilities of both investigated salts were corrected for the hydrolysis

reaction and correlated via a semi-empirical approach based on the phase

equilibrium between the present phases.

Keywords: Calcium chloride, Magnesium chloride, Solubility, Supercriti-

cal water

- 110 -

5.1 ∥ Introduction

5.1 Introduction

The design of industrial processes on supercritical fluids highly depend on the quality

and accuracy of the property data that are - if at all - available for the relevant systems.

While a broad range of systems of supercritical CO2 plus organic compounds of all kinds

have been investigated due to advantages that came along with the usage of supercritical

CO2, this is not the case for most other fluids (e.g. Propane, Methanol). The same

applies for supercritical water.

Despite the corrosiveness and mechanical stress that supercritical water represents to

equipment and material at elevated elevated temperature and pressure (Tc = 647K, pc =22.1MPa), it has been considered as a medium of choice for reactions, polymerization,

destruction of waste components, gasification of biomass and particle formation (1–5).

Yet, the limited amount of property data on systems consisting of supercritical water

and organic / inorganic compounds is noted (6).

In many systems undergoing supercritical water processing, salts and other inorganic

compounds are present to some degree. Examples for these systems are waste streams in

supercritical water oxidation, biomass and other fuels in supercritical water gasification,

and impurities in feed water streams (7; 8). Since water in its supercritical state loses

its polar character and thereby its ability to dissolve inorganic compounds in more than

minimal quantities, salts precipitate and start to form a solid phase. The presence of

such an additional phase can have a major influence on the process and cause unwanted

and unrecognized side effects in the process itself. Such operations can be effected in the

long term by corrosion and erosion of the equipment. The salts can also act as catalysts

(e.g. alkali salts in the water-gas shift reaction during the gasification of biomass (9))

and avoid coke and tar formation (e.g. coke and tar formation in gasification processes

(9)). Another possibility of the formation of an additional phase is the option to remove

this phase from the system (e.g. by gravity, by centrifugal forces) and thereby separate

the present salts from the remaining water (10).

To analyze the behavior of inorganic compounds and to enlarge the available property

data base, the authors have investigated the solubility of mono-valent alkali nitrates

(LiNO3, NaNO3, KNO3) and alkali chlorides (LiCl, NaCl, KCl) (11; 12). The focus

of this work is on the solubility of MgCl2 and CaCl2 to extend the available data to

bivalent salts and to continue the systematic investigation of solubilities in supercritical

water. These salts have been investigated in the range of 660 to 690 K and 18.5 to

23.5 MPa. Furthermore, the formation of hydroxides will be disccused, which took

- 111 -

Chapter 5 ∥ Magnesium Chloride and Calcium Chloride

place during the experiments. These results were correlated with an approach based on

a phase equilibrium between the present phases and compared to the previous works.

5.2 Experimental

The measurements of the solubilities was performed using a continous flow method. The

experimental setup and method has been described in detail elsewhere (11; 12); thus

only the most significant information are presented in the following.

The scheme of the experimental apparatus is shown in Figure 5.1, permitting mea-

surements up to 723 K and 25 MPa. Hastelloy is the material of choice for all heated

parts. The pressure in the system was established via a HPLC pump (LabAlliance Series

III, LabAlliance, USA), while a custom-made oven provided heat. An U-tube is installed

inside the oven with a length of 265 mm, an inner diameter of 4.6 mm and an outer

diameter of 6.35 mm. The temperature in the oven was measured at the inlet, at a

middle position and at the outlet via standard Type K thermocouples.

HPLC pump

Supply vessel

Preheater

Cooling

Back Pressure

Regulator Relief Valve

TI-1

Filter

2 m

Oven

Salt

column

Preheater

temperature

Oven Inlet

temperature

Temperature control

oven

Outlet

temperature

Pressure

Analysis

temperature

Analysis

and

samplingConduc-

tivity

measurement

Oven Outlet

temperature

Oven Center

temperature

TI-2

TI-3

TI-5 CI-1

PI-1

TI-4

TI-6

TC-1

Figure 5.1 ∥ Scheme of the experimental setup

Upon entering the U-tube, the feed stream can become supersaturated depending

on the temperature, pressure and feed concentration. If an oversaturation occurs, the

excess amount of salt will precipitate until the phase equilibrium between both phases

is established in the column. The exiting stream leaves the system in equilibrium and at

the solubility resulting from the temperature and pressure in the column. The stream

- 112 -

5.2 ∥ Experimental

is cooled down and depressurized to ambient conditions. Samples are taken when an

equilibrium state is verified via measurement of the conductivity of the outlet stream.

The analysis of the samples is done via an inductive coupled plasma atom emission spec-

trometer (ICP, Perkin-Elmer Optima 5300DV, Perkin-Elmer, USA, uncertainty < 2 %

for all investigated species) for the concentrations of Mg and Ca, ionic chromatography

(IC, Metrohm 741 Compact IC, Metrohm AG, Switzerland, uncertainty < 5 % for all

investigated species) was used for the chlorine concentration. The pH of all samples

was measured after the actual experimental run with a standard pH electrode (WTW

pH/Cond 340i/SET, WTW Wissenschaftlich-Technische Werkstatten GmbH, Germany,

uncertainty after calibration ± 0.01). For the calculation of the density, the outlet tem-

perature (TI-4 in Figure 5.1) and the pressure at the pressure sensor (PI-1 in Figure 5.1)

were used. The feed solution was prepared with deionized water and analytical grade

MgCl2 and CaCl2 (Boom B.V., The Netherlands).

5.2.1 Analysis of the precipitates

If precipitates were found in the column, samples of these precipitates were taken from

the inside of the U-tube. For optical analysis, the specimens of the solid material were air

dried, mounted on specimen studs, sputtered with a thin gold layer by using a sputtering

unit (Jeol JFC-1200, Jeol, Japan) and imaged with a scanning electron microscope

(Jeol JSM-6480LV, Jeol, Japan) at 6 kV . The SEM was combined with a energy

disperve X-ray spectroscopy unit (Noran System SIX, Thermo Fisher Scientific, USA).

In addition, images were taken via a standard light microscope (Leica MZ9.5 / DFC 320,

Leica Microsystems, Germany). Parts of these precipitates were found to have formed

bigger, loose particles in the bottom of the U-tube (cf. Figure 5.2). Other parts of the

precipitates were found on the walls of the tubing forming small clusters. Figure 5.2

shows a cross section of the U-tube where precipitates can be found attached to the wall

of the tubing. The clusters could easily be scrapped off the wall.

The typical EDX distribution for the precipitates can be found in Table 5.1. As can

be seen from the table, the samples consist mainly of magnesium and oxygen and only

contain traces of chloride. Other metals and carbon in the distribution result from the

surrounding tubing material which is made of Hastelloy.

Further analysis of the precipitates was performed via infrared spectrography (ATR-

FTIR, Model Shimadzu 8400S, measurement range between 400 to 4000 cm−1). The

result of the IR analysis of the sample can be found in Figure 5.3. As can be seen,

- 113 -

Chapter 5 ∥ Magnesium Chloride and Calcium Chloride

(a)

(b)

Figure 5.2 ∥ Precipitates from the bottom of the U-tube; image taken via light micro-

scope (a); Precipitates on the column wall; image taken via SEM (b)

one peak at about 3690 cm−1 can be clearly distinguished. This peak results from the

hydroxyl group of magnesium hydroxide (13; 14).

Combining the EDX and the IR results, it is to conclude that the formed particles

consist of magnesium hydroxide. This conclusion is also supported by the shape of the

crystals. Magnesium hydroxide tends to form hexagonal crystals which could also be

found in the particles found here (cf. Figure 5.4) (15). Particles consisting of MgCl2

can not be found at ambient state due to the high solubility in water at these conditions.

- 114 -

5.3 ∥ Results and Discussion

Table 5.1 ∥ EDX distribution

Element wt.-%

Carbon 23.00

Oxygen 43.83

Magnesium 23.06

Chlorine 0.34

Chrome 0.08

Manganese 0.01

Nickel 9.69

Figure 5.3 ∥ IR spectra of the precipitates

5.3 Results and Discussion

5.3.1 Correlation of the experimental data

For the interpretation of the experimental results, a description on base of a phase equi-

librium is chosen (16–19). Here, it is assumed that the solid phase and the supercritical

fluid form an equilibrium depending on the state of the system (cf. Figure 5.5). The

equilibrium between the solid and fluid phase can be formulated as follows:

- 115 -

Chapter 5 ∥ Magnesium Chloride and Calcium Chloride

Figure 5.4 ∥ Image of the hexagonal crystals of Mg(OH)2

a ⋅Mec ∗ m ⋅H2O(f) + b ⋅Xd ∗ p ⋅H2O(f)⇌ MeaXb ∗ n ⋅H2O(f) (5.1)

MeaXb ∗ n ⋅H2O(f) ⇌ MeaXb(s) + n ⋅H2O(f) (5.2)

Ô⇒Ks =α

MeaXb ∗ n⋅H2O(f)

αMeaXb(s)

⋅ αnH2O(f)

(5.3)

Here, Me and X represent the salt cation respectively the salt anion with a and b as the

number of ions in the salt molecule and c and d their valency. s and f refer to the phases

solid and fluid; n, m and p are the number of water molecules. The formation of the

solid phase is assumed to occur via the associated complex and not via dissociated ions.

The phase equilibrium constant Ks on base of the activities of the present species can

be simplified with several assumptions. The interaction between the presented species

is neglected while the activity coefficient of the solid salts is assumed as unity. The fluid

phase is assumed as an ideal one, thereby allowing the usage of the density of water as

concentration. More elaborated description of the assumptions and the efficiency of this

approach can be found be elsewhere (11; 16; 17). The assumptions lead to the following

expression for the solubility of an inorganic compound in supercritical water:

K∗s ≈

mMeaXb ∗ n⋅H2O(f)

1 ⋅ ρnm, H2O(f)

(5.4)

Ô⇒ mMeaXb ∗ n⋅H2O(f)

=K∗s ⋅ ρnm, H2O(f)

(5.5)

- 116 -

5.3 ∥ Results and Discussion

Figure 5.5 ∥ Equilibrium curve of NaCl in supercritical water (11)

This expression can be extended further by substituting the equilibrium constant under

usage of a van’t Hoff-like expression:

K∗s ≈

mMeaXb ∗ n⋅H2O(f)

1 ⋅ ρnm, H2O(f)

(5.6)

Ô⇒ mMeaXb ∗ n⋅H2O(f)

=K∗s ⋅ ρnm, H2O(f)

(5.7)

Ô⇒ log mMeaXb ∗ n⋅H2O

= logK∗s + n ⋅ log ρ

m, H2O(5.8)

= −∆solvH

R ⋅ T + ∆solvS

R+ n ⋅ log ρ

m, H2O(5.9)

R is the universal gas constant, T the system temperature, m the solution molality, ρ

the density. Ks respectively K∗s are the equilibrium constant and the equilibrium constant

including the simplifications. The Gibbs energy of solvation, ∆solvG, the enthalpy of

solvation, ∆solvH, and the entropy of solvation, ∆solvS are assumed as independent of

the system parameters temperature, pressure and density; n is later on referred to as

the coordination number. The molalities and densities are expressed on an amount of

substance base. The density of pure water is calculated via the IAPWS95 equation of

state (20). The experimental data were fitted to the parameters ∆solvH, ∆solvS and n.

More information on this approach can be found in the previous works of the authors

(11; 12).

- 117 -

Chapter 5 ∥ Magnesium Chloride and Calcium Chloride

5.3.2 Experimental results on Magnesium Chloride

The solubility of MgCl2 was investigated in the range of 660 to 670 K and 19 to 22.5

MPa.

Contrary to previous experiments (11; 12), precipitates were found inside the column

after rinsing and cleaning the setup. These precipitates are assumed to result from

a parallel hydrolysis reaction. Additionally, a drop in pH from appr. 6.7 for the feed

solution to between 2.5 and 2.9 depending on temperature and pressure was recorded

(cf. Table C.1).

Hydrolysis mechanism

The mechanism that is assumed to lead to the formation of the precipitates and to be

responsible for the pH decrease, is depicted below:

MgCl2 + 2 H2O ⇋ Mg(OH)2(s) ↓ +2 HCl (5.10)

(5.11)

The formed magnesium hydroxide precipitates and remains in the column while the

formed HCl remains in solution and leaves the U-tube, whereby the pH of the effluent

stream decreases. This is supported by the fact that all investigated samples had a

pH between 2.4 and 2.9 while the feed solution had a pH of 6.8. The occurrence of a

hydrolysis reaction was also found by other authors (17; 21).

Solubility of Magnesium Chloride

Resulting from the composition of MgCl2, the ratio between the magnesium and chlorine

concentration in the samples should be two. In Figure 5.6, the measured anion and cation

solubilities can be seen. While the anion concentration shows a steadily increasing trend

with increasing density, the cation concentration shows deviations from the expected

behavior as could be assumed after former investigations of other salts. Also, the ratio

between both concentrations is not two, but higher and does not have a constant value.

It is assumed that these deviations result from the parallel hydrolysis reaction.

In order to correct the measured solubilities for this parallel reaction, the following

correction for the magnesium concentration can be made via the formulation of a com-

ponent balance, the via the ICP measured magnesium concentrations and pH values and

Eq. 5.11:

- 118 -

5.3 ∥ Results and Discussion

Figure 5.6 ∥ Uncorrected results of the magnesium and chlorine composition as a

function of water density; ○, chlorine concentration; ▽, magnesium concentration, △,

double magnesium concentration

cH3O+ = cHCl = 2 ⋅ cMg(OH)2 (5.12)

ctotal(Mg) = cICP (Mg) + cHydrolysis(Mg)= cICP (Mg) + c(Mg(OH)2)= cICP (Mg) + 0.5 ⋅ c(HCl)= cICP (Mg) + 0.5 ⋅ 10−pH (5.13)

Figure 5.7 shows the original magnesium and chlorine concentration, the corrected

magnesium concentration and - in order to compare the ratios between the chlorine con-

centration and the corrected magnesium concentration - twice the corrected magnesium

concentration. As can be seen, the corrected magnesium concentration now shows a

consistent trend. Furthermore, the ratio between the corrected magnesium concentra-

tion and the chlorine concentration is approximately two for the samples presented here.

Therefore, a correction of the magnesium concentration for the hydrolysis reaction is

assumed as necessary to evaluate the experimental data.

The experimental results of MgCl2 including the corrections can be found in Figure

5.8. In order to correlate the experimental data with the equilibrium approach mentioned

above (cf. Eq. 6.3), the parameters ∆solvH, ∆solvS and n were fitted via a minimization

- 119 -

Chapter 5 ∥ Magnesium Chloride and Calcium Chloride

Figure 5.7 ∥ Corrected results of the magnesium and chlorine composition as a function

of water density; ○, chlorine concentration; ◻, magnesium concentration; △, corrected

magnesium concentration; ▽, double magnesium concentration

routine included in MATLAB to the experimental data. As can be seen, the experimental

data is in good agreement with the correlation.

The values for the three parameters ∆H, ∆S and n can be found in Table 5.2.

The experimental data including temperature and pressure as well as their standard

deviations, the density, the composition and the pH for the measurements on MgCl2

can be found in Table C.1.

5.3.3 Experimental results on Calcium Chloride

The solubility of CaCl2 was investigated in the range of 660 to 690 K and 18.5 to 23.5

MPa.

During the experiments with CaCl2, no precipitates were found after rinsing and cleaning

the setup in contrast to the experiments with MgCl2. The non-presence of any precipi-

tates of calcium compounds is assumed to result from the different solubility products as

discussed below. A drop in pH however from appr. 6.7 for the feed solution to between

3.4 and 4 was found indicating nevertheless the occurence of the parallel hydrolysis re-

action (cf. Tbl. C.2 and Fig. 5.9). Therefore, the experimental data of CaCl2 was

corrected in the same manner as described above for MgCl2.

- 120 -

5.3 ∥ Results and Discussion

Figure 5.8 ∥ Solubility of MgCl2 as a function of water density; ○, this work; solid line

represents the description of the experimental data with Eq. 6.3

Figure 5.9 ∥ Measured pH of the CaCl2 samples as a function of water density; ▽,

measured pH values

The presence of precipitates at ambient state after rinsing

Crystals of MgCl2 and CaCl2 were not found after the experiments due to their higher

solubilities at ambient conditions. For the presence of the products of the hydrolysis

- 121 -

Chapter 5 ∥ Magnesium Chloride and Calcium Chloride

Table 5.2 ∥Model parameters of the salts MgCl2 and CaCl2 for Eq. 5.9

Salt ∆H/J ⋅mol−1 ∆S/J ⋅mol−1 ⋅K−1 n / - Molecule radius (22) / 10−12m

MgCl2 2 -107.42 3.44 419

CaCl2 8436 -82.57 2.52 468

reaction, the solubility product of Mg(OH)2 and Ca(OH)2 at 298 K, 0.1 MPa can be

compared. Ca(OH)2 has a solubility product of 5.02 ⋅ 10−6, while the solubility product

of Mg(OH)2 is with 5.61 ⋅ 10−12, several orders of magnitude smaller (15). Therefore,

no particles of Ca(OH)2 can be found while Mg(OH)2 is still present at ambient state

after rinsing with water.

Solubility of Calcium chloride

The experimental results of CaCl2 including the corrections can be found in Figure 5.10.

As can be seen, the experimental data is in good agreement with the correlation.

The values for the three parameters ∆H, ∆S and n can be found in Table 5.2. The

experimental data including temperature and pressure as well as their respective standard

deviations, the density, the composition and the pH for the measurements on CaCl2 can

be found in Table C.2.

5.3.4 Possible correlation between salt properties and model pa-

rameters

As described in the previous work of the authors (12), a correlation between the parame-

ters derived from Eq. 5.9 and the radius of the molecules is investigated. The radii of the

molecules is calculated as the sum of the corresponding crystal radii. The parameters for

three alkali nitrates (LiNO3, NaNO3, KNO3) respectively chlorides (LiCl, NaCl, KCl)

plus the parameters of three additional monovalent salts (CuO,PbO,KOH) combined

with the parameters presented in the section before are used for investigating this pos-

sible correlation. The parameters for CuO,PbO and KOH were derived from property

data available in literature (19; 23; 24). The parameters for these salts and the alkali

chloride and nitrate salts can be found in Table 5.3 (12). The radii of the salt molecules

were obtained from the corresponding crystal radii and can be found in Table 5.3 (22).

The figures 5.11, 5.12 and 5.13 contain the respective parameter of the salts as a

- 122 -

5.3 ∥ Results and Discussion

Figure 5.10 ∥ Solubility of CaCl2 as a function of density; ○, this work; dashed line

represents the description of the experimental data with Eq. 6.3

function of the radius. Two classes of salts can be distinguished for the parameter n

depending on their bonding character (12). The results for ∆H and ∆S do not allow

such a clear conclusion due to experimental errors and the limited range of experimental

data. Yet, the authors assume a comparable correlation between these parameters and

the property of the respective salt molecule.

Table 5.3 ∥Model parameters of the salts CuO, PbO, and KOH for Eq. 5.9 (12)

Salt ∆H/J ⋅mol−1 ∆S/J ⋅mol−1 ⋅K−1 n / - Molecule radius (22) / 10−12m

CuO 23749 -103.68 1.34 238

PbO 27901 -57.97 1.97 197

KOH 13369 -81.10 3.24 277

LiCl 5199 -69.61 2.48 240

NaCl 18784 -86.74 4.88 280

KCl 13123 -95.66 4.65 318

LiNO3 15594 -81.91 4.33 352

NaNO3 5149 -93.06 3.93 392

KNO3 -7793 -111.91 3.72 430

- 123 -

Chapter 5 ∥ Magnesium Chloride and Calcium Chloride

Figure 5.11 ∥∆H (cf. Eq. 5.9) as a function of the molecule radius; ○, CuO; ▽, PbO;

◇, LiCl; ◻, KOH; △, NaCl; 7, KCl; ⊕, LiNO3; ⋆, NaNO3; +, MgCl2; ⊞, KNO3; |,

CaCl2

Figure 5.12 ∥∆S (cf. Eq. 5.9) as a function of the molecule radius; ○, CuO; ▽, PbO;

◇, LiCl; ◻, KOH; △, NaCl; 7, KCl; ⊕, LiNO3; ⋆, NaNO3; +, MgCl2; ⊞, KNO3; |,

CaCl2

- 124 -

5.4 ∥ Conclusions

Figure 5.13 ∥ n (cf. Eq. 5.9) as a function of the molecule radius; ○, CuO; ▽, PbO;

◇, LiCl; ◻, KOH; △, NaCl; 7, KCl; ⊕, LiNO3; ⋆, NaNO3; +, MgCl2; ⊞, KNO3; |,

CaCl2

5.4 Conclusions

In the work presented here, the solubilities of MgCl2 and CaCl2 were studied in de-

pendence of the parameters density, temperature and pressure. The investigated range

was 660 to 690 K and 18.5 to 23.5 MPa. The measurements were performed using a

continuous flow method.

For all experiments, a decrease in pH was found that was caused by a parallel hydrol-

ysis reaction. An approach to correct this parallel reaction was presented and applied

successfully to interpret the experimental results. For the experiments with MgCl2, pre-

cipitates were found after rinsing the setup in contrary to any former experiments. These

precipitates were to found to be consisting of Mg(OH)2, thereby proving the occurrence

of the parallel hydrolysis reaction. The reason for the presence of these precipitates at

ambient state after rinsing the setup in comparison to any other salt investigated so far

is assumed to result from the low solubility product of Mg(OH)2.

The corrected experimental results could be correlated in good agreement with Eq.

5.9. The parameters derived from this correlation agree to the trends for these param-

eters presented in the earlier studies of the authors, where a dependency between the

parameters and the radius of the salt molecule was found (12).

- 125 -

Chapter 5 ∥ Magnesium Chloride and Calcium Chloride

Further it is to conclude that parallel reactions like presented in this work have to be

kept in mind for further investigations and applications. Although this might be consid-

ered as a minor problem for certain systems with less severe hydrolysis and thereby lower

pH variations like NaCl, it can lead to errors in measurement and evaluation of solubil-

ities if not addressed properly. Also, the presence of precipitates even at ambient state

must be taken into account in order to avoid damage to the equipment and disturbance

of measurements and system behavior.

Acknowledgements

The authors would like to thank Thibaut Garcia de Changy for his contribution to the

experimental part of this work. Additionally the authors would like to thank Kamuran

Yasadi, Arie Zwijnenburg, Janneke Tempel and Jelmer Dijkstra for their contribution in

the analysis of the samples.

This work was performed in the TTIW-cooperation framework of Wetsus, centre of

excellence for sustainable water technology (www.wetsus.nl). Wetsus is funded by the

Dutch Ministry of Economic Affairs, the European Union Regional Development Fund,

the Province of Fryslan, the City of Leeuwarden, and the EZ/Kompas program of the

’Samenwerkingsverband Noord-Nederland’. The authors like to thank the participants

of the research theme Salt for their financial support.

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[3] M. D. Bermejo, M. J. Cocero, Supercritical water oxidation: A technical review, AIChE Journal52 (11) (2006) 3933–3951.

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- 126 -

5.5 ∥ References

[8] Y. Matsumura, T. Minowa, B. Potic, S. R. A. Kersten, W. Prins, W. P. M. van Swaaij, B. van deBeld, D. C. Elliott, G. G. Neuenschwander, A. Kruse, M. J. Antal, Biomass gasification in near-and super-critical water: Status and prospects, Biomass & Bioenergy 29 (4) (2005) 269–292,0961-9534.

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[10] I. Leusbrock, S. J. Metz, G. Rexwinkel, G. F. Versteeg, The removal of inorganic compoundsfrom water streams via supercritical water, Proceedings of the 9th International Symposium onSupercritical Fluids.

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