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Removal of inorganic compounds via supercritical waterLeusbrock, Ingo

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Citation for published version (APA):Leusbrock, I. (2011). Removal of inorganic compounds via supercritical water: fundamentals andapplications. Rijksuniversiteit Groningen.

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This chapter has been published as:Leusbrock, I., Metz, S. J., Rexwinkel, G., and Versteeg, G. F.; The solubilities of phos-phate and sulfate salts in supercritical water ; The Journal of Supercritical Fluids 54(1),1-8.

Chapter 6

The solubilities of phosphate andsulfate salts in supercritical water

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Chapter 6 ∥ Phosphate and sulfate salts

Abstract

Inorganic compounds are regularly present in aqueous streams. To understand

their influence and behavior on these streams at supercritical conditions, little to

no property data is available, which can be used as starting point for further

research or application design. Since inorganic compounds tend to precipitate

at these conditions, scaling, blocking and erosion can occur as a consequence.

Furthermore, a separation of (precious) compounds from the bulk stream due

to the precipitation is possible. Here, phosphate compounds are regarded as

interesting for further investigation since resources are assumed to be depleted

in future. As phosphate is present in many waste streams, these could be used

as sources for recoverable phosphate. Resulting from these facts and options,

a proper understanding and knowledge of these systems is important for later

industrial applications. Therefore, the authors have investigated the behavior of

salts (e.g. NaCl, NaNO3 and MgCl2) in supercritical water in previous works.

To extend this knowledge, the solubilities of the sulfate salts MgSO4 and

CaSO4 in a range of 18.8 - 23.2 MPa and 655 - 675 K as well as of the phosphate

salts Na2HPO4, NaH2PO4 and CaHPO4 in a range of 20.5 - 24.2 MPa and

665 - 690 K were investigated in this work with a continuous flow method in

continuation of former work of the authors. The solubilities were compared with

existing data available from open literature. A quantitative correlation on base of

a phase equilibrium between the present phases was used to describe the behavior

and to compare it with previous results. For the investigated calcium salts, CaSO4

and CaHPO4, it was found that a significant solubility decrease already happens

at subcritical conditions resulting in precipitation in unwanted locations. For the

remaining compounds, a parallel hydrolysis reaction was found as could be seen

from a change in pH in the effluent stream.

Keywords: Phosphates, Solubility, Sulfates, Supercritical water

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6.1 ∥ Introduction

6.1 Introduction

Supercritical water (SCW) is regarded as one of the most promising and interesting

solvents / environments for future industrial applications such as hydrothermal conver-

sion of biomass (1; 2), destruction of waste sludge and organic waste compounds (3; 4),

particle formation on a micro- and nanoscale (5; 6), reaction media for polymerizations

and conversion processes (7; 8) and separation of inorganic compounds. Due to funda-

mental changes in properties at supercritical conditions (critical temperature Tc = 647

K, critical pressure pc = 22.1 MPa), SCW varies to a large degree in its properties

compared to the common phases liquid and vapor. The additional benefit of supercriti-

cal fluids and SCW is the fact that properties like diffusion rate, density and viscosity are

adjustable in the supercritical regime (9; 10). As a result of the comparably harsh criti-

cal conditions of water1 and a limited number of industrial and commercial installations,

available property data on systems containing supercritical water is scarce in comparison

to other supercritical fluids.

Inorganic compounds can be found in a majority of aqueous streams. At supercrit-

ical conditions, these compounds tend to precipitate or to form vapor or liquid phases

saturated in salt. This precipitation can lead to scaling and plugging in installations and

therefore has to be controlled or avoided (11; 12). First investigated in connection with

supercritical water oxidation (SCWO) (13; 14), it became obvious in later applications

like supercritical water gasification (SCWG) that the presence of inorganic compounds

can have a major influence on the reaction and the process itself (e.g. as a catalyst

(15)). Furthermore, severe problems can occur for a continuous long-term process (e.g.

corrosion (16; 17), increased pressure drops due to scaling (12), catalyst poisoning (18)).

Therefore, concepts were developed to avoid the formation of such a solid phase (e.g.

different reactor modifications (19; 20)) or to separate inorganic compounds in an inte-

grated treatment step before the actual process step (18; 21). For the design of these

new concepts and even more sophisticated ones like a fractionation of different salt

fractions, it is vital to have a sound understanding of the solubility of these inorganic

compounds and their influence and behavior. One of the group of compounds, which are

of interest for a separation from other compounds, are phosphates. Phosphates are part

of many waste streams, which represents a potential source for recoverable phosphate

(22). As phosphate resources are limited and a shortage or even a depletion of global

1Carbon dioxide as the most frequently applied supercritical fluid has a critical temperature Tc =

304 K and a critical pressure pc = 7.4 MPa.

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resources is expected, new ways have to be found to recover phosphate and to guarantee

sufficient fertilizers for global food production.

In previous works of the authors, monovalent salts (alkali nitrates and chlorides (23;

24)) and bivalent salts (calcium and magnesium chloride (25)) were investigated. Here,

a continuous flow method was used to investigate the solubilities of these compounds.

In these studies, different quantitative approaches to describe solubilities in SCW were

compared. The behavior and a quantitative description for the solubility of monovalent

chloride and nitrate salts and double-valent chloride salts was presented. In the course

of this study, the solubilities and behavior of the bivalent salts MgSO4 and CaSO4 in a

range of 18.8 - 23.2 MPa and 655 - 675 K as well as of the mono- / bivalent cation

/ triple-valent anion phosphate salts Na2HPO4, NaH2PO4 and CaHPO4 in a range of

20.5 - 24.2 MPa and 665 - 690 K was object of the investigations. The experimental

data was correlated with an approach derived in an earlier work (23). The results of

this correlation are compared with previous results of the authors as well as with data of

similar compounds available in literature (26).

6.2 Theoretical background

6.2.1 Correlation of the experimental data

For quantitative description of the solubilities, an approach on base of a phase equilibrium

between the present phases is applied (23; 27). In this approach, the solid phase and the

supercritical fluid are assumed to form an equilibrium depending on the conditions in the

system (23). Thereby, a description of the equilibrium between the phases is possible in

the following manner:

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

Ô⇒Ks =α

MeaXb ∗ n⋅H2O(f)

αMeaXb(s)

⋅ αnH2O(f)

(6.1)

Me and X resemble the salt cation / the salt anion; a and b is the number of ions in

the salt molecule; c and d their valency; s and f refer to the solid and fluid phase; n, m

and p to the number of water molecules. For this approach it is furthermore assumed

that a formation of the solid phase occurs via the associated complex and not directly

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6.3 ∥ Experimental

via dissociated ions (28). For the equilibrium constant Ks, several simplifications can be

made. Interaction between the species in the system is neglected; the activity coefficient

of the solid phase is assumed as unity. The fluid phase is treated as an ideal one. A

more extensive description of this approach can be found be elsewhere (23; 24; 27). As a

result of these assumptions, the solubility of inorganic compounds in supercritical water

can be described in the following way:

K∗s ≈m


1 ⋅ ρnm, H2O(f)

Ô⇒ mMeaXb ∗ n⋅H2O(f)

=K∗s ⋅ ρnm, H2O(f) (6.2)

Substituting the equilibrium constant under usage of a van‘t Hoff-like expression leads

to the following expression:

K∗s ≈m


1 ⋅ ρnm, H2O(f)

Ô⇒ mMeaXb ∗ n⋅H2O(f)

=K∗s ⋅ ρnm, H2O(f)Ô⇒ log m

MeaXb ∗ n⋅H2O= logK∗s + n ⋅ log ρm, H2O

= −∆solvHR ⋅ T +

∆solvS

R+ n ⋅ log ρ

m, H2O(6.3)

For a more convenient interpretation of the parameters, a description on an amount of

substance base is chosen. Ks and K∗s are the equilibrium constant and the equilibrium

constant including the simplifications. R is the universal gas constant, T the system

temperature, m the molality, ρ the density, n the coordination number. 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. The density of pure water is calculated via the IAPWS95 equation

of state (29). Experimental data was fitted to the parameters ∆solvH, ∆solvS and n to

quantitatively describe the solubility. A more detailed description on this approach can

be found elsewhere (23; 24).

6.3 Experimental

For the measurements of the solubilities of the phosphate and sulfate salts, a continuous

flow method was applied like in the previous works of the authors. The most important

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details of the method and setup are presented in the following; a more detailed description

can be found elsewhere (23; 24).

The setup is designed for an operation up to 25 MPa and 723 K. A scheme of the

setup can be found in Figure 6.1. Heat was provided by a custom-made oven and a

pre-heater; for pressurization and flow, a HPLC Pump was used (LabAlliance Series III,

LabAlliance, USA). An U-tube was 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

inside the oven was measured at inlet, middle and outlet position by standard Type K

thermocouples (relative uncertainty 0.25 %). The pressure inside the setup is measured

with a pressure sensor (Keller PA23H, relative uncertainty 0.2 %). The material of choice

for all heated parts was Hastelloy, the diameter of every tubing was 1/16” (=1.588 mm)

if not mentioned otherwise.

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 6.1 ∥ Scheme of the experimental setup

The feed stream enters the U-tube, where an oversaturation of the feed stream de-

pending on the system conditions (pressure, temperature, density) and the feed con-

centration can occur. If so, the exceeding amount starts to precipitate and form an

additional phase till an equilibrium in the system has established. The stream exits the

U-tube at an equilibrium state with a composition resulting from the system conditions.

After cooling and depressurization, samples can be taken. The analysis of these samples

is done via an inductive coupled plasma atom emission spectrometer (ICP, Perkin-Elmer

Optima 5300DV, Perkin-Elmer, USA, uncertainty < 2 % for all investigated species) for

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6.4 ∥ Results and Discussion

the concentrations of Na, Ca, Mg and P. Ionic chromatography (IC, Metrohm 741 Com-

pact IC, Metrohm AG, Switzerland, uncertainty < 5 % for all investigated species) wasused for measurement of the SO4 and PO4 concentration. Conductivity measurement of

the outlet stream was used for the verification of an equilibrium state in the column. pH

measurements of the samples were performed using an standard pH electrode (WTW

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

uncertainty after calibration ± 0.01). The outlet temperature (TI-4 in Fig. 6.1) andthe pressure at the pressure sensor (PI-1 in Fig. 6.1) were used for the calculation of

the density in the system. The feed solution was prepared with deionized water and

analytical grade compounds (Boom B.V., The Netherlands).

6.4 Results and Discussion

6.4.1 Experimental results phosphate salts

The solubility of Na2HPO4, NaH2PO4 and CaHPO4 was investigated in a range of 20.5

- 24.2 MPa and 665 - 690 K. For the sodium phosphates a feed concentration of 0.05 M

was used; for CaHPO4 a feed concentration of 0.05 - 0.5 mM was used2. Na2HPO4 and

NaH2PO4 showed occasional plugging, apparently independent of temperature, pressure

and length of the experiment. The plugging in this case occurred at the filter in the outlet

section, which could be removed by rinsing with nitric acid (cf. Fig. 6.1). CaHPO4 on

the other hand showed severe plugging in the preheater and the inlet section. Although

the preheater was switched off to avoid plugging in later experiments with CaHPO4,

plugging of the preheater and the inlet section occurred before an equilibrium in the

column was achieved. Therefore, no further experiments on CaHPO4 were performed.

Since CaHPO4 has a low solubility at ambient state (∼200 mg ⋅ L−1 (30)), a furtherdecrease to minimal solubilities at elevated temperatures was to be expected resulting in

the observed plugging. For all experiments with phosphate salts, the composition of the

samples was analyzed in regard of the content of Na+, P and PO3−4 . Each data point

results from the experimental results of four to six individual samples at one temperature

and pressure.

Phosphates form a complex conjugated acid / base system of the acid / base pairs of

H3PO4, H2PO−4 , HPO2−4 and PO

3−4 (cf. Fig. 6.2), wherein all forms act as weak acids

/ bases and as buffer (also see Eq. 6.4 to 6.6).

2The feed concentration was decreased in order to avoid plugging

- 135 -


Figure 6.2 ∥ Relative fraction of phosphate species as a function of pH

Figure 6.3 ∥ Solubility of Na2HPO4 as a function of density; ○, this work; dashed linerepresents the description of the experimental data with Eq. 6.3

H3PO4 +H2O ⇋ H2PO−4 +H3O+ (pK = 2.12) (6.4)H2PO

−4 +H2O ⇋ HPO2−4 +H3O+ (pK = 7.21) (6.5)

HPO2−4 +H2O ⇋ PO3−4 +H3O+ (pK = 12.63) (6.6)

- 136 -


By changes of the pH in the system, the equilibria and the relative fractions of the

different steps can be changed.

(a)

(b)

Figure 6.4 ∥ Solubility of sodium and phosphate (a) and phosphorus and phosphate(b) as function of density for the experiments on Na2HPO4; ○ corresponds to sodiumconcentration; ◻ corresponds to phosphorus concentration; ▽ corresponds to phosphateconcentration

Disodium hydrogen phosphate

Figure 6.4 shows the concentrations of the three analyzed compounds (Na+, P and PO3−4 )

as a function of the molar density. Each data point consists of five samples of one ex-

- 137 -


perimental run at one set temperature and pressure. The pH value for these samples

were in the range of 6.5 - 7.5, while the pH of the 0.05 M feed solution was calculated

as 9.18 (calculated with OLITM , a simulation software for electrolyte chemistry). As can

be seen, the concentration of phosphorus (analyzed by ICP) and phosphate (analyzed by

IC) coincide within experimental errors. It can be concluded that no decomposition of

the phosphate happened. According to the molar composition of Na2HPO4, a factor 2

between the concentration of phosphate / phosphorus to sodium was to be expected in

the samples. Regarding the sodium concentration, one can see that all measured con-

centrations (with one exception) were only slightly above the corresponding phosphorus

/ phosphate concentration and did not reach a ratio of 2.

A possible explanation for this could be a parallel hydrolysis like described in the

following equation:

Na2HPO4 +H2O ⇋ NaOH ↓ +NaH2PO4 (6.7)

The hereby produced sodium hydroxide is assumed to have a lower solubility than

the other present compounds and to precipitate at the experimental conditions. The

effluent stream consequently would leave the system at a lower sodium / phosphate

ratio than the influent stream. A continued reaction to phosphoric acid can be excluded

since the effluent pH was between 6.5 and 7.5. Resulting from this, only H2PO4−

and HPO42− are present in the system (cf. Fig. 6.2). A correction of the cation

concentration like proposed in Eq. 6.9 does not lead a significant increase in the cation

concentration. It is assumed that the buffer capacity of the system prevents a further

decrease in pH. Thereby, a direct correlation of the pH to the amount of precipitated

NaOH cannot be derived. For the further quantitative description of the solubility, the

sodium concentration is used in accordance to previous works. One has to be aware that

the actual solubility of Na2HPO4 is underestimated to a certain degree by this approach.

The sodium concentration for the experiments with Na2HPO4 can be found in Figure

6.3. Furthermore included in this graph is the description of the experimental data with

Eq. 6.3. As can be seen, the approach is in good agreement with the experimental data.

More details on the approach and fitting procedure can be found in section 6.2 and in

previous works (23; 24).

The actual experimental data containing temperature, pressure, pH and sodium con-

centration can be found in Table D.1; the parameters for Eq. 6.3 for Na2HPO4 can be

found in Table 6.1.

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Sodium dihydrogen phosphate

Figure 6.5 shows the concentrations of Na+ and PO3−4 as a function of the molar density.

Each data point consists of five samples of one experimental run at one set temperature

and pressure. The data for the phosphorus concentration coincides to a high degree

with the phosphate concentration and is therefore not included in this figure. The pH

value for these samples were in the range of 2.7 - 3.2, while the pH of the 0.05 M

feed solution was calculated as 4.54 (calculated with OLITM , a simulation software for

electrolyte chemistry). According to the molar composition of NaH2PO4, the concentra-

tions of phosphorus / phosphate and sodium were expected to be equal. Yet, all sodium

concentrations were below the corresponding phosphorus / phosphate concentrations.

A hydrolysis reaction is assumed as the reason for this in accordance with previous ob-

servations (25). This hydrolysis leads to a formation of NaOH, which precipitates. The

parallel formed HCl stays in solution and leaves the system with the effluent. This leads

to a decrease in pH as can be found for these experiments. A correction of this hydrolysis

reaction is possible with the following approach:

mH3O+ = mH2PO−4 =mNaOHmtotal(Na) = manalyzed(Na) +mHydrolysis(Na)

= manalyzed(Na) +m(NaOH)= manalyzed(Na) + 10−pH (6.8)

As can be seen from Figure 6.5, the correction shows in this case a good coincidence

with the results for the phosphate / phosphorus results. For further evaluation of the

results, the corrected sodium concentration is used. The corrected sodium concentration

for the experiments with NaH2PO4 can be found in Figure 6.6. Furthermore included in

this graph is the description of the experimental data with Eq. 6.3. As can be seen, the

approach is in good agreement with the experimental data. More details on the approach

and fitting procedure can be found in section 6.2 and in previous works (23; 24).

The actual experimental data containing temperature, pressure, pH and corrected

sodium concentration can be found in Table D.2; the parameters for Eq. 6.3 for

Na2HPO4 can be found in Table 6.1.

- 139 -


(a)

(b)

Figure 6.5 ∥ Solubility of sodium and phosphate (a) and comparison between the cor-rected solubility of sodium and the solubility of phosphate (b) as function of density for

the experiments on NaH2PO4; ○ corresponds to sodium concentration; ▽ correspondsto phosphate concentration; △ corresponds to corrected sodium concentration

Comparison phosphate salts

Figure 6.7 shows a comparison of the results for the phosphate salts Na2HPO4 and

NaH2PO4 presented in this work and the results for KH2PO4 by Wofford et al. (26).

The comparison shows that the solubility of potassium phosphate in supercritical water

is lower than the solubility for sodium phosphates. This agrees with the assumption that

the molecule size or in this case the cation size is related to the solubility as presented

- 140 -


Figure 6.6 ∥ Solubility of NaH2PO4 as a function of density; ○, this work; dashed linerepresents the description of the experimental data with Eq. 6.3

Figure 6.7 ∥ Comparison of the solubilities of Na2HPO4, NaH2PO4 (this work) andKH2PO4 (26) as function of density; dashed line corresponds to Na2HPO4; solid line

corresponds to NaH2PO4; dashed-dotted line corresponds to KH2PO4

in previous works (24; 31). In these works, a direct relation between the cation of the

investigated chloride salts and the actual solubility was found (Li > Na > K). As ions

- 141 -


are present in supercritical water only to a small amount, the majority of the non-water

species present is associated as molecules and complexes. To keep these molecules in

solution, hydration has to occur. This hydration is due to the decreased amount and

strength of hydrogen bondings in supercritical water less present than at ambient state.

Therefore, a hydration of larger molecules is more difficult as more hydrogen bondings

have to be established over a larger distance (24).

Upon a comparison between the solubilities of KH2PO4 and the two sodium phos-

phates, it can be seen that the solubility of KH2PO4 is lower in agreement with the

findings for the alkali chlorides and nitrates. As one compares the solubilities of the

phosphates investigated in this work, NaH2PO4 has a slightly higher solubility than

Na2HPO4 in the investigated density range. It can be concluded here as well that the

molecule size is of influence since Na2HPO4 and the resulting hydration sphere is larger

by a small amount than NaH2PO4 and thereby resulting in the lower solubility.

6.4.2 Experimental results sulfate salts

Two sulfate salts that are commonly found in water streams, MgSO4 and CaSO4, were

investigated in the range of 18.8 - 23.2 MPa and 655 - 675 K. CaSO4 has a low

solubility product at ambient conditions (Ksp = 4.93 ⋅ 10−5 (30)), while the solubilityof MgSO4 is comparably high (Ksp = 4.67 (30)). Therefore, a low feed concentration

(1 mM and below) was used to avoid plugging. Despite this low feed concentration,

plugging occurred in and just before the preheater for both salts. The tubing in the

preheater has a diameter of 1.588 mm and thereby only a small cross section. A further

reduction in concentration did not solve this issue and just postponed the point of

plugging by several minutes to hours. To solve this problem, the preheater was turned

off in order to avoid precipitation before the solution enters the oven. This was successful

for MgSO4, yet not for CaSO4. Due to thermal conductance, parts of the preheater

and parts just after the preheater were heated up apparently far enough to precipitate

parts of the feed solution in case of CaSO4. This indicates a higher solubility of MgSO4

at supercritical pressures and slightly elevated temperatures in comparison to CaSO4

similar to the behavior of CaHPO4 (cf. previous section). Therefore, only results for

the measurements of MgSO4 can be presented here. Attempts to remove the plugging

via ultrasound and rinsing with concentrated nitric acid / Piranha solution (3:1 conc.

H2SO4 to 30 % H2O2) did not succeed for both salts.

- 142 -


Magnesium sulfate

The experimental data for the measurements of MgSO4 can be found in Figure 6.8.

The overall solubilities are low and in the range of 10−4 to 10−5 mol ⋅kg−1. Thereby, theconcentrations were in the lower detection range of the analytical methods applied here.

Furthermore included in this graph is the quantitative description of the experimental

data via Eq. 6.3.

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

all samples, a pH value of between 3.9 and 4.6 was measured, while the feed solution of

1 mM MgSO4 a pH value of 6.98 was determined (calculated with OLITM , a simulation

software for electrolyte chemistry). This decrease in pH indicates the occurrence of a

parallel hydrolysis reaction (25). Nevertheless, a correction of the cation concentration

for a parallel hydrolysis reaction as done for MgCl2 and CaCl2 was not performed due

to the following. The values for the cation and anion concentration coincide and the

deviation between the anion and cation concentration was within the experimental error

(cf. Fig. 6.9). Therefore, no excess amount of one of the species magnesium or sulfate

was present in the sample which would make a correction necessary. If one performs

a correction of the cation concentration for the hydrolysis as proposed in Leusbrock et

al. (25) (also see Eq. 6.9), a difference of up to 80 % between the anion and the

cation concentration (cf. Fig. 6.9) would be the result. This was not considered as a

reasonable approach for the evaluation of the experimental results. Still, one has to be

aware that hydrolysis occurs as indicated by the pH decrease.

mH3O+ = 2 ⋅mH2SO4= 2 ⋅ mMg(OH)2

mtotal(Mg) = manalyzed(Mg) +mHydrolysis(Mg)= manalyzed(Mg) +m(Mg(OH)2)= manalyzed(Mg) + 0.5 ⋅ m(H2SO4)= manalyzed(Mg) + 0.5 ⋅ 10−pH (6.9)

The actual experimental data containing temperature, pressure, pH and magnesium

concentration can be found in Table D.3; the parameters for Eq. 6.3 for MgSO4 can be

found in Table 6.1. In contrast to the results on the phosphate salts, each data point

represents one sample. Thereby, the accuracy of the measurement is lower.

- 143 -


Figure 6.8 ∥ Solubility of MgSO4 as a function of density; ○, this work; dashed linerepresents the description of the experimental data with Eq. 6.3

Table 6.1 ∥ Model parameters of the salts MgSO4, Na2HPO4 and NaH2PO4 for Eq.6.3

Salt ∆H/J ⋅mol−1 ∆S/J ⋅mol−1 ⋅K−1 n / -MgSO4 -8257 -154.4 3.31

Na2HPO4 -142411 -351.4 5.37

NaH2PO4 -35582 -152.4 3.47

KH2PO4 -84569 -271.2 4.33

6.5 Conclusions

In the course of this work, the solubilities of the sulfate salts MgSO4 and CaSO4 in a

range of 18.8 - 23.2 MPa and 655 - 675 K as well as of the phosphate salts Na2HPO4,

NaH2PO4 and CaHPO4 in a range of 20.5 - 24.2 MPa and 665 - 690 K were investi-

gated. It has been found that experiments with CaSO4 and CaHPO4 lead to plugging

of the equipment with the current configuration of the setup due to low solubilities even

at a subcritical state. Therefore, no solubility data can be presented at this point for

these salts.

For the other salts, experiments have been conducted successfully. Here, deviations

- 144 -

6.5 ∥ Conclusions

(a)

(b)

Figure 6.9 ∥ Solubility of magnesium and sulfate (a) and comparison of sulfate concen-tration and possible correction (b) as function of density for the experiments of MgSO4;

○ corresponds to sulfate concentration; △ corresponds to magnesium concentration; ◻corresponds to corrected magnesium concentration

between the anion and cation concentrations occurred due to a parallel hydrolysis re-

action. These could partly be corrected by inclusion of the measured pH values. The

extent of the hydrolysis depends on the temperature and pressure during the experiments

and the investigated compound as found in previous articles of the authors (24; 25). The

results of the experiments were correlated with a semi-empirical model. This correlation

showed a good coincidence between experiments and proposed model.

- 145 -


For a better understanding of the behavior of phosphate salts in supercritical water,

more experiments with further salts and different conditions have to be conducted due

to the complexity (pH dependency, equilibria between different stages) of the phosphate

system.

6.6 Acknowledgements

The authors would like to thank the analytical team of Wetsus 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 Fryslân, 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.

6.7 References

[1] S. Stucki, F. Vogel, C. Ludwig, A. G. Haiduc, M. Brandenberger, Catalytic gasification of algaein supercritical water for biofuel production and carbon capture, Energy & Environmental Science2 (5) (2009) 535–541.

[2] G. Brunner, Near critical and supercritical water. Part I. Hydrolytic and hydrothermal processes,The Journal of Supercritical Fluids 47 (3) (2009) 373.

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