A biphase H₂O/CO₂ system as a versatile reaction medium for organic synthesis

Abstract

We review activity on the usage of a biphase H₂O/CO₂ system in organic synthesis as a reaction medium of green chemistry. The formation of self-neutralizing carbonic acid in such a system eliminates the problem of salt disposal, typical for acid-catalyzed reactions with the usual mineral and organic acids. A large variety of different reactions to be performed in the biphase H₂O/CO₂ system are discussed in detail, including cyclization/cycloaddition, hydroformylation, hydrogenation, reduction, coupling, rearrangement, substitution, addition, halogenation, hydrolysis, oxidation and others. These reactions cover a significant part of modern organic synthesis. The main physical properties of carbonic acid being formed in the biphase H₂O/CO₂ system and their dependence on the temperature and pressure of saturating CO₂ are analyzed. The problem with the search for the most optimal reaction conditions from the viewpoint of selection of appropriate pressure and temperature regions for the best yields and selectivity achievable is addressed in general. Comparison with formation and utilization of peroxycarbonic acids, alkylcarbonic acids and carbamic acids by means of saturation with pressurized CO₂ of some other biphase systems is discussed in relationship to organic synthesis as well. The influence of a CO₂ admixture on the unique properties of high temperature water, another promising green solvent, is also considered.

---

Marina A. Pigaleva

Marina A. Pigaleva joined associate professor Marat O. Gallyamov's research group at Lomonosov Moscow State University as a student in 2008. She graduated with honors from Physics Faculty of Moscow State University in 2013. After that she proceeded her research in the same laboratory as a PhD student. From 2014 she is working as a researcher and high-pressure media specialist in the Professor Aziz M. Muzafarov's research group in the Nesmeyanov Institute of Organoelement Compounds. Her recent investigations mainly focus on the features of organization of cationic polymers in the solutions of carbonic acid at high pressure of saturating CO2.

Igor V. Elmanovich

Igor V. Elmanovich graduated from M.V. Lomonosov Moscow State University, Faculty of Physics in 2011. He received PhD in Polymer Science from A.N. Nesmeyanov Institute of Organoelement Compounds in 2015. His main research interest lies in implementation of compressed or supercritical CO2 in various applications. Among these applications are: organosilicon chemistry in carbonic acid solutions, electrocatalytic materials preparation using supercritical CO2 as a medium and deposition of water- and oil-repellent coatings on fabrics and nonwoven materials from solutions in supercritical CO2.

Yuriy N. Kononevich

Yuriy N. Kononevich graduated from Chernihiv national Pedagogical University named after T. H. Shevchenko, Ukraine. He performed his PhD study at the Institute of Pharmacology and Toxicology of the National Academy of Medical Sciences of Ukraine. He is currently working at the Institute of Organoelement Compounds of Russian Academy of Sciences. His research interests are synthesis of organometallic compounds, organoboron dyes, organic luminescent materials and photochemistry.

Marat O. Gallyamov Marat O. Gallyamov

Marat O. Gallyamov received his PhD degree in 1999 from Lomonosov Moscow State University (Russia). In 2002–2004 he stayed as Humboldt Research Fellow with Prof. Dr Martin Möller at Universität Ulm (Germany). In 2009 he received DSc degree (habilitation) from Lomonosov MSU. Since that time he is leading the “Laboratory of Polymers at Surfaces and New Materials for Fuel Cells” (Lomonosov MSU) as an Associate Professor. In 2012 he received Shuvalov Award (Lomonosov MSU). One of his current research interests is focused on application of compressed carbon dioxide systems for design and synthesis of new materials for advanced applications.

Aziz M. Muzafarov

Aziz M. Muzafarov is a Full Member of Russian Academy of Sciences since 2011 and a Director of Nesmeyanov Institute of Organoelement Compounds of RAS since 2013. He received his PhD degree in Polymer Science from Nesmeyanov Institute of Organoelement Compounds of RAS, Russia (1981). Since 1991 Aziz Muzafarov is a Head of Laboratory at Enikolopov Institute of Synthetic Polymer Materials of RAS, Russia. He received DSc degree (habilitation) in Chemistry from Karpov Institute of Physical Chemistry, Russia (1997). His main research interests involve organosilicon compounds, structure/properties relationship, supramolecular organization, hybrid organic–inorganic structures, chlorine-free silanes technology, functional metallosiloxanes, dendrimers.

---

1. Introduction

The biphase H₂O/CO₂ system has attracted significant attention recently as a versatile reaction medium for different types of synthetic reactions. In this system both phases are condensed: liquid water is in contact with pressurized carbon dioxide, which has therefore a relatively high density. In the majority of cases described in the literature (with yet some exceptions), the density of the liquefied or supercritical (temperature > 31 °C) carbon dioxide is quite comparable to the density of water. Interesting, that the both phases are benign from the environmental viewpoint, which is an unusual situation for biphase systems typically applied in organic synthesis nowadays, such as water/oil systems. Therefore, the organic synthesis to be performed in the biphase H₂O/CO₂ system follows the paradigm of green chemistry. In the biphase H₂O/CO₂ system water is, naturally, a polar phase, whereas CO₂ is a non-polar “oil-like” phase. Indeed, CO₂ has zero dipole moment, though rather strong quadrupole moment, which may bring on some residual polarity as well as certain polarizability. Dissolving power of compressed CO₂ is known to depend strongly on pressure. Reagents, catalysts (if used), intermediates, and products may have different solubility in the two phases, thus their distribution between them may be tuned by the pressure variation. That introduces an additional flexibility as well as a convenient tool to optimize selectivity and reaction pathways in general.

In distinct from typical biphase systems, cross-contaminations of the two phases in the biphase H₂O/CO₂ system is not an issue.¹ Indeed, H₂O “contaminates” CO₂ phase rather poorly, as it will be discussed below. Yet, slightly humidified CO₂ phase is still a strongly non-polar fluid, which may be dried if necessary in the reaction course rather easily. Further, CO₂ may “contaminate” H₂O phase more substantially. This “contamination” results in formation of solutions of carbonic acid with the pH value of about 3 and somewhat below at high pressures, as it will be discussed. Yet, aqueous solutions of carbonic acid are still quite polar media, pH values of which may be tuned in certain range by the pressure of the saturating CO₂ phase. After decompression acidity of this phase is decreased due to release of dissolved CO₂. The pH values are accordingly increased up to the range of 4–5.5, and by means of bubbling with some inert gas they may be increased even further up to quite neutral values of about 7. This easily neutralized by a decompression and subsequent gas bubbling carbonic acid may serve as a convenient and environmentally friendly substituent for typical mineral and organic acids currently used in organic synthesis as reaction media, which require neutralization with alkaline and thus inevitably demand salt disposal.

Working with pressurized biphase H₂O/CO₂ system generally implies a necessity to use proper high pressure setup. From the experimental standpoint this is a certain limitation. Nevertheless, the industrial and laboratory high pressure equipment was significantly improved during the last decades. Indeed, firstly, many implemented and quite established industrial process of chemical engineering anyway require rather high pressures. Besides, secondly, the general interest towards supercritical (sc) fluids as pressurized green solvents was ever-increasing recently.^2–11 As a result, new easy-to-use and powerful pressure generators as well as sophisticated and convenient autoclaves are nowadays widely available for research laboratories and industry in general. Possibility of automation of continuous organic reactions performed in the presence of pressurized CO₂ was successfully demonstrated.¹²

It is interesting to mention that the presence of CO₂ may affect the course of certain reactions. Indeed, CO₂ may participate directly in some of them, for example it is quite reactive with epoxides, thus it may also serve as C1 synthon.¹ Further, if reduction reaction is performed, CO₂ may be reduced to CO, which is capable to poison catalyst surface (such as surface of Pt).¹ This possibility should be taken into account. Fortunately, CO₂ is completely oxidized, therefore it cannot participate directly in oxidation reactions, in distinct from many other non-polar solvents.¹

Our own recent experimental experience has demonstrated that the biphase H₂O/CO₂ system is particularly useful in silane chemistry.¹³ Here the possibility to vary the pH value by means of physical parameters (pressure, temperature) serves as a key factor determining the composition of the end products and, moreover, allowing to control and to tailor it. The results¹³ demonstrated that it is very promising to extend this experience onto other new classes of organic and organoelement compounds.

In the literature there was only one previous review on organic reactions in biphase H₂O/CO₂ system, but in Chinese language (published in 2010).¹⁴ To the best of our knowledge, it was never translated in English. Yet, this review¹⁴ may be quite useful as far as it contains relevant references (by that time) and structures of corresponding reactions. There were also some other reviews focused on broader class of gas-expanded liquids in general, where biphase H₂O/CO₂ system (as a reaction medium and catalytically active one) was mentioned as well, though rather briefly without any detail focus on its unique physical and chemical properties.^15–18 At the same time, many original papers were recently published on this particular topic. The majority of citations to be discussed and analyzed in our review has been never generalized and summarized before thus still requiring detailed and comprehensive analysis. Therefore, we believe that the present review is strongly demanded.

Our paper appeals to chemical engineers and chemical scientists, particularly those, who are working with sc CO₂. The interest towards sc CO₂ is growing no more as far as this system is already quite well studied and understood. Yet, the convenient equipment for sc CO₂ applications is widely available nowadays. It is quite suitable for experiments and technological processes involving biphase H₂O/CO₂ system, carbonic acid, etc. Focusing on this topic the chemical community may learn how it is possible to replace in a variety of established chemical processes the usual acids with self-neutralizing carbonic acid (meaning ultimate solution of salt disposal problem in acid catalysis).

Thus in this review we analyze peculiar physical properties of the biphase H₂O/CO₂ system as well as describe main directions of organic synthesis currently being performed in this promising reaction medium at normal and elevated temperatures. The paper is organized as follows. In the next section, main physical properties of biphase H₂O/CO₂ system will be described and discussed. Then the following three sections will summarize relevant works on organic synthesis in the presence of carbonic acid (biphase H₂O/CO₂ system) as well as will provide comparison with the reactions to be performed in the presence of peroxycarbonic acid, alkylcarbonic acids, carbamic acids. In the last review section the influence of high temperature on the organic synthesis will be discussed. Finally, the paper is concluded, relevant comments on proper reaction conditions are given and promising research directions are highlighted.

2. Physical properties of biphase H₂O/CO₂ system

2.1 Formation of carbonic acid in biphase H₂O/CO₂ system

It is well known that water in a contact with carbon dioxide becomes acidic due to formation and dissociation of carbonic acid. In the solutions of carbonic acid a complex equilibrium exists, which can be described in general as follows:

If water is saturated with carbon dioxide at high pressure, this equilibrium shifts towards the formation and dissociation of carbonic acid, therefore increasing the acidity of the medium. Indeed, the thus formed molecules of carbonic acid dissociate and generate protons, which allow to decrease the pH down to the value of 3 and somewhat less at CO₂ saturating pressure values of about 30 MPa (see discussion below for further details). High pressure setup including a reaction vessel (see Fig. 1) is required for the formation of carbonic acid. It should be noted that working with high pressures may be dangerous and thus requires usage of certified equipment as well as the compliance with proper safety regulations!

Fig. 1 The scheme of the main pathway of the reversible reaction for CO₂ after its dissolution in water with the formation of molecules of carbonic acid and their subsequent dissociation into protons and bicarbonate (mainly) ions.

When CO₂ is dissolved in water (Fig. 1), three chemical reactions occur and six reagents participate in them: water (H₂O), carbon dioxide (CO₂), carbonic acid (H₂CO₃), protons (H⁺), bicarbonate ions (HCO₃⁻) and carbonate ions (CO₃²⁻). The corresponding chemical reactions for the biphasic water/CO₂ medium are the assembling of H₂CO₃ molecules through the interaction of dissolved CO₂ molecules and dissociating H₂O molecules, as well as the first and second dissociations of carbonic acid molecules. Those processes are typically described by the following set of equations, where the dissolution of CO₂ in the water medium foregoes the reactions:¹⁹

Recently MD simulations²⁰ of the biphase H₂O/CO₂ system showed that the dissolved CO₂ (CO_2(aq) in reactions (2) and (3)) is a poor hydrogen bond acceptor and is extremely weakly hydrated. The radial distribution functions for water containing dissolved CO₂ indicate a stronger interaction between the CO₂ carbon and the water oxygens, which is a prerequisite for the hydrolysis of dissolved CO₂ to form carbonic acid, bicarbonate, and carbonate ions (reactions (3)–(5)). A soft X-ray absorption spectroscopy performed by the same group earlier²¹ showed in turn that carbonic acid (H₂CO₃ in reactions (3) and (4)) is able to donate two strong hydrogen bonds to solvating waters while acting as a weak acceptor of one such bond.

The solubility constant K₀ for the process (2) can be expressed from the Henry law:

where [CO_2(aq)] is the concentration of CO₂ dissolved in water, P_CO2 is the CO₂ partial pressure.

The dissociation constants (K′, K₁ and K₂) for the reactions (3)–(5) can be written as follows:²²

Using these relationships it is easy to demonstrate that the value of K′ can be expressed by the following equation:

K′ = k_H2CO3/k_CO2 = (K₁/K^ap₁) − 1, (10)

where k_CO2 and k_H2CO3 are the rate constants of the forward and backward reactions (3), respectively.

The rate constant of the backward reaction (3) is much higher than the rate constant of the forward reaction (3). Consequently, the concentration of the dissolved CO₂ in water in equilibrium is much higher (99.8%) than the concentration of H₂CO₃ (0.2%).²³ The concentration of bicarbonate ion in its turn is comparable to that of the carbonic acid and the concentration of CO₃²⁻ species is reported to be negligible.²⁴

The apparent dissociation constant commonly measured in the experiments (0.1 MPa CO₂, 25 °C):²⁵

is about 600 times smaller (0.1 MPa CO₂, 25 °C)²⁵ than the true dissociation constant K₁ (8). The reason for that is the fact that using common techniques one cannot reliably distinguish between dissolved CO₂ and H₂CO₃ molecules, particularly at so small relative concentration of the latter ones. Therefore, usually, one measures only the integral content of the molecules of both types. As a result, this experimental limitation misleadingly provides the decreased apparent value of K^ap₁ (11) instead of expectedly higher real value of K₁ (8).

Using eqn (10) the true dissociation constant K₁ (0.1 MPa CO₂, 25 °C) can be expressed and approximated by the following equation, considering that K′ ≈ 600:

K₁ = K^ap₁ × (1 + K′) ≈ 2.5 × 10⁻⁴ (12)

It was found out²⁶ that the value of the apparent first ionization constant decreases with temperature increase (when the temperature increases from 25 °C to 250 °C at the CO₂ pressure of 0.1 MPa the K^ap₁ decreases by 25 times, see Fig. 2) and increases with pressure increase (more than 7-fold increase in the K^ap₁ value when the CO₂ pressure is increased from 0.1 MPa to 200 MPa at 25 °C, see Fig. 3).

Fig. 2 Apparent first ionization constant of carbonic acid as a function of temperature at pressure equal to 0.1 MPa (data points were taken from ref. 26).

Fig. 3 Apparent first ionization constant of carbonic acid as a function of pressure at temperature equal to 25 °C (298 K) (data points were taken from ref. 26).

Recently the true pK_a value for carbonic acid was directly measured by means of the infrared active marker modes of a photoacid (2-naphthol-6,8-disulfonate).²⁷ The experimental data obtained by this method confirmed that for the carbonic acid at atmospheric CO₂ pressure and room temperature pK_a = 3.45 ± 0.15, as opposed to the apparent pK_a of 6.35, when carbonic acid is defined as both dissolved CO₂ in water and actually formed carbonic acid H₂CO₃. Therefore it can be concluded that carbonic acid is a stronger acid then it was previously believed. The acidity of carbonic acid is comparable, for example, to that of formic acid (pK_a ≈ 3.8).

2.2 Solubility of CO₂ in H₂O in biphase H₂O/CO₂ system

The investigation of the solubility of CO₂ in water originated from the pioneer works of Henry,²⁸ where he postulated his famous solubility law. The solubility of CO₂ in water at low and moderate pressures is well described by this law (see eqn (6)). The deviations from Henry law begin at pressures around 10 MPa and higher. Therefore it has been known for a long time that the solubility of CO₂ in water monotonically increases with the increase of pressure.

From the very useful and detailed report of Spycher et al.²⁹ one can conclude that a typical graph of the solubility of CO₂ in water with the increase of pressure can be divided into two main parts, which differ in the solubility slopes (see Fig. 4). The first part of the graph with the higher slope of the solubility/pressure dependence is a part where CO₂ is in a state of a low-density gas.

Fig. 4 Solubility of CO₂ in water as a function of pressure at (a) 25 °C, (b) 60 °C, (c) 110 °C. Reprinted from ref. 29, copyright© 2003, with permission from Elsevier.

As it is well known, CO₂ is in a state of gas up to the pressure of liquefaction if the temperature is below critical one. It should be noted that low temperatures promote the liquefaction of gases. Indeed, the attraction forces between molecules should be strong enough for their “sticking together” to form a liquid. Increased pressure can act as an assistant in this as far as it results in a reduction of intermolecular distance on overall, which imply stronger van der Waals attraction between molecules. But higher temperatures lead to the activation of molecular thermal motion that counteracts the intermolecular attraction. The higher the temperature, the higher pressure is required to liquefy a gas. Moreover above certain temperature (critical temperature, equal to 31 °C for CO₂) a gas cannot be liquefied at all. Above the critical temperature and critical pressure (7.39 MPa for CO₂) any gas including CO₂ is supposed to be in a supercritical state.

The second part of the graphs (Fig. 4) with the somewhat smaller slope of the solubility/pressure dependence corresponds either to liquefied CO₂ region if the temperature is below the critical temperature, or to dense supercritical CO₂ region when the temperature is above 31 °C.

Thus, the solubility of the gaseous CO₂ in water is proportional to the partial pressure according to the Henry law. Whereas in the region of liquid or dense supercritical state, the solubility of CO₂ in water depends less pronouncedly on the pressure, mainly because the molar volume of liquid CO₂ varies to a lesser extent with the pressure than the molar volume of gaseous CO₂. But it should be mentioned, that the dependence of the molar volume of liquid CO₂ on the pressure is nevertheless more pronounced, than the same parameter for water, which in the context of this aspect can be considered as a distinctly incompressible liquid.

Considering the dependence of CO₂ solubility in water on the temperature it is very important first of all to emphasized its non-monotonic behavior. This peculiar behavior determines the dependence of efficiency of many organic reactions to be performed in biphase H₂O/CO₂ system on the reaction temperature. The fact that the solubility of CO₂ in water is a non-monotonic function of the temperature can be seen already from comparison of Fig. 4a–c. If one compares the graphs it becomes obvious that there is some minimum of the solubility at certain intermediate temperatures. The existence of the minimum of the solubility of CO₂ in water correlates also with weakened acidity of the medium at the correspondingly elevated temperatures. But it should be noted that the pH of the carbonic acid will increase with the temperature also because the ability of H₂CO₃ to dissociate decreases with the temperature rise, see the discussion of properties of high temperature water below.

The non-monotonic behavior of the solubility of CO₂ in water with the temperature increase³⁰ can be more clearly seen in Fig. 5. Typically, there is a certain minimum of the solubility achieved at intermediate temperatures, the localization of which yet depends on the pressure. The higher the pressure, the lower is the position of the minimum at the temperature scale. For example, if the pressure is about 70 MPa, the minimum is localized at 70 °C, approximately.

Fig. 5 Solubility of CO₂ in water: non-monotonic dependence on temperature at different pressures, reprinted from ref. 30, copyright© 1941, with permission from American Chemical Society.

2.3 Solubility of H₂O in CO₂ in biphase H₂O/CO₂ system

In the region of gaseous CO₂ the solubility of water in CO₂ decreases with the pressure rise at the fixed temperature. This is a typical behavior for solubility of liquids (i.e., their vapors) in compressed gases.³¹ Indeed, if the temperature is fixed, the partial pressure of a vapor equilibrated with its liquid is a constant value. Therefore, the total pressure may increase only due to the increase of concentration (density) of the compressed gas. In a simplest assumption of ideal mixing, the molar concentration of the vapor in its mixture with the pressurized gas should be directly proportional to the ratio of the saturated vapor pressure at this temperature to the total pressure, i.e. the sum pressure of partial contributions of both the vapor and the compressed gas. As far as the former is a constant, that means the concentration of the vapor in its mixture with the compressed gas should be inversely proportional to the total pressure (hyperbolic law). Of course, this is a very simplified picture and many corrections are to be taken into account in reality, e.g. Poynting effect. The Poynting effect implies the following. If the compressed gas exerts pressure on the liquid, this should affect also the vapor, which is in equilibrium with that liquid. As a result, the pressure of the vapor should be also somewhat higher than the saturated vapor pressure at this temperature in the absence of the effect of the compressed gas. Nevertheless, with the possible relevant corrections taken into account, the solubility of water in compressed gases is still expected to decrease monotonically with the pressure increase and to follow approximately the law, similar to the simplest hyperbolic decline. This behavior was observed experimentally, for example, by Bartlett for the dependence of solubility of water in some “indifferent” gases, such as hydrogen or nitrogen at normal and elevated temperatures.³¹ Bartlett used the word “indifferent” gases for hydrogen and nitrogen meaning that they cannot be liquefied at the experimental temperatures as far as their critical temperatures are located very significantly lower. At these conditions it is also impossible to bring their densities closer to the typical liquid densities at experimentally achievable pressures. On the contrary, CO₂ may be liquefied if the experimental temperature is below 31 °C at pressures of around several MPa. Thus it naturally demonstrates quite different behavior for the solubility of water in it with the pressure rise. Moreover, even if CO₂ is transferred into a sc state (T > 31 °C) at moderately elevated temperatures, it can be compressed to the densities of about 1 kg l⁻¹ (that is typical for liquids) when experimentally quite achievable pressure (several dozens of MPa) is applied.

When the transition of CO₂ being pressurized from a gaseous to a liquid state occurs, there is a typical sharp jump on the solubility curves that is followed by the change of the type of the dependence: from the sharp decrease to the smooth increase with the pressure rise. This jump can be simply explained by the following considerations: while more and more gaseous CO₂ being pressurized enters in the closed vessel with constant volume, some amount of the carbon dioxide starts to liquefy in order to free some space for the newly arriving gas molecules, but the pressure of the system does not change. The subsequent increase in pressure only can occur when all the CO₂ in the vessel turns into liquid. Indeed, according to the van der Waals isotherm (see Fig. 6) under quasi-static compression, starting from the point B, the system splits into two phases, i.e. a liquid and a gas. The density of the liquid and gas remains unchanged during the compression, and equal to their values at the points of A and B, respectively. The pressure remains constant. The amount of substance in a gas state decreases continuously, while the amount of substance in a liquid state increases until the point A is reached. In this point all the substance has been transformed into the liquid state. The localization of the jump on the solubility curves (Fig. 7) particularly corresponds to the liquefaction pressure. The uncertainty of solubility at this point, where the solubility can take on any values from a corresponding range, correlates directly to the volume (density) uncertainty of the liquefied gas in the range AB (Fig. 6), i.e. at the liquefaction pressure. Thus, we have a transition from solubility of H₂O in a gas CO₂ to solubility of H₂O in a liquid CO₂.

Fig. 6 Typical van der Waals isotherm for a real (non-ideal) gas.

Fig. 7 Solubility of H₂O in carbon dioxide as a function of pressure at (a) 25 °C, (b) 60 °C, (c) 110 °C. Reprinted from ref. 29, copyright© 2003, with permission from Elsevier.

In the region of liquid CO₂ the solubility of water in CO₂ phase increases slightly with pressure rise, see Fig. 7. The solubility of water in CO₂ phase also increases with temperature (Fig. 7). For comparison, quasi-hyperbolic law valid for an ideal mixing of a vapor with a pressurized “indifferent” (in a Bartlett's meaning³¹) gas is presented in Fig. 7 as a dot line.

So, when the temperature is below the critical one for CO₂, there is the typical jump on the dependence of solubility of H₂O in carbon dioxide on pressure, as discussed above, see Fig. 7a. This jump is localized at the point of transition of CO₂ from a gas to a liquid state (the pressure of liquefaction). The position of this transition point depends on temperature: it shifts towards higher pressures with the temperature rise. But, if the temperature is above the critical one (31 °C for CO₂) the sharp jump disappears. Instead of the jump there remains to be only a local minimum of the solubility of H₂O in carbon dioxide and also a point of inflection, where the solubility curve becomes convex instead of being concave. Similarly to the transition between gaseous and liquid CO₂, marked by the jump, this minimum may be considered as a border line between a “supercritical gas” state and a “supercritical liquid” state. In both cases, this is a supercritical fluid as far as the temperature is above the critical one and the pressure is still higher than the critical one (31 °C and 7.4 MPa for CO₂). But if the density of the sc fluid is relatively low (to the left of the minimum on the pressure scale), solubility of water in it decreases with pressure, as it does for solubility of water in compressed “indifferent” gases (see above). When its density becomes relatively high (to the right of the minimum on the pressure scale), solubility of water in it increases with pressure, as it does for solubility of water in liquids (see above). Therefore, it is only natural to consider this minimum of the solubility of H₂O in CO₂ as a border between “supercritical gaseous” CO₂ (lower density) and “supercritical liquid” CO₂ (higher density) regions. It is interesting that the localization of this border shifts towards higher pressures (10 MPa and above) with temperature increase, see Fig. 7.

Thus, above the critical temperature of CO₂, the sharp jump (Fig. 7a) on the H₂O solubility curve with the increase of pressure becomes smoother (Fig. 7b and c). Indeed, the transition from a “supercritical gas” state to a “supercritical liquid” state occurs continuously, not abruptly, with continuous density increase, without formation of any phase boundary, in distinct from the situation with compressed gas being liquefied.

The total amount of water to be dissolved in CO₂ phase in comparison with the typical CO₂ solubility in water is approximately by an order of magnitude smaller (compare Fig. 5 and 7). Teng et al.³² explained this feature by the difference in properties of molecular carbon dioxide and water. Indeed, carbon dioxide has non-polar molecules with a zero dipole moment and with mainly London dispersion forces acting between them. Whereas water has strongly polar molecules with a dipole moment equal to 6 × 10⁻³⁰ C × m and with intermolecular interactions determined significantly by the hydrogen bonds. In the biphasic mixture of CO₂ and H₂O molecular interactions between two equal molecules (H₂O and H₂O, or CO₂ and CO₂) are much stronger than between two different ones (H₂O and CO₂). But electrostatic forces of H₂O molecules can polarize initially nonpolar CO₂ molecules, thereby increasing their ability to penetrate into the polar water phase. Whereas CO₂ molecules cannot force the water molecules to become less polar ones. Consequently, the CO₂ solubility in water is much higher than the H₂O solubility in CO₂ (Fig. 5 and 7).

2.4 pH of aqueous phase in biphase H₂O/CO₂ system

As it was mentioned before, the pH value of the water phase in the biphase H₂O/CO₂ system is determined by the amount of carbonic acid formed, when CO₂ is dissolving in water. While the CO₂ saturating pressure increases, more and more CO₂ molecules dissolve in water, thus more H₂CO₃ molecules are formed and dissociate, mainly into HCO₃⁻ and H⁺. Consequently, the pH dependence on the pressure shows that the pH values monotonically decrease with the increase in pressure at the moderate temperatures (see Fig. 8a and 9).^33–35 The increase of the temperature correspondingly leads to the decrease of the solubility of CO₂ (Fig. 5) in water and to the decrease of the acidity constant thus the pH values increase with temperature rise (see Fig. 8b and 9).^33,34

Fig. 8 pH values of biphase H₂O/CO₂ system as a function of (a) pressure: , T = 35 °C; , T = 50 °C; , T = 70 °C; , T = 95 °C; , T = 125 °C; T = 150 °C. (b) Temperature: , p = 0.38 MPa; , p = 0.61 MPa; , p = 1.00 MPa; , p = 2.43 MPa; , p = 6.23 MPa; , p = 9.28 MPa; , p = 15.4 MPa. Reprinted from ref. 33, copyright© 2013, with permission from Elsevier.

Fig. 9 Variation of pH of water in equilibrium with saturating pressurized CO₂ at temperatures of (♦) 25 °C, (■) 40 °C, (▲) 50 °C, and (●) 70 °C with pressure rise. Reprinted from ref. 34, copyright© 1995, with permission from American Chemical Society.

It is interesting to mention, that amines may become protonated at low pH values. Therefore, some polymers containing amino groups become polycations and thus dissolve in aqueous phase of the biphase H₂O/CO₂ system. For example, it was shown that biopolymer chitosan may be dissolved in the pressurized aqueous solutions of carbonic acid and moreover chitosan macromolecules self-organize into peculiar elongated rod-like nanoaggregates in this medium.^36–38 It was also possible to construct micelles from polystyrene-block-poly(4-vinylpyridine) block copolymer in the pressurized carbonic acid solutions due to protonation of poly(4-vinylpyridine) and plasticizing action of compressed CO₂ towards the polymers.³⁹

2.5 Densities of both phases in biphase H₂O/CO₂ system

According to King et al.⁴⁰ there is no measurable difference in the density of the CO₂ phase in biphase H₂O/CO₂ system as compared to the density of pure CO₂ at the same pressures. This is due to the low solubility of water in CO₂ phase, which is discussed above. Of course, the density of the CO₂ increases with pressure rise.

On the contrary, the water phase density changes noticeably when water phase is in contact with pressurized CO₂ phase: the higher the pressure, the higher the mole fraction of the dissolved CO₂ in water and therefore the more pronounced is the change in water density.^41–43

The water density increases faster with pressure when water is in contact with gaseous CO₂, then when it is in contact with liquid CO₂. This feature can be explained by the principles of the CO₂ solubility in water that were discussed above. The curves in Fig. 10 undergo a sharp bend in the point of CO₂ transition from a gas to a liquid state at the temperatures below critical one. But they demonstrate more gradual and smooth evolution without so sharp bend above this temperature (when supercritical gaseous CO₂ transforms gradually to supercritical liquid CO₂).

Fig. 10 CO₂-saturated water density as a function of pressure and temperature. The temperature ranges from 11 °C to 60 °C. Reprinted from ref. 41, copyright© 2004, with permission from American Chemical Society.

As reported by Hebach et al.⁴¹ the largest differences between the densities of pure water and water saturated with CO₂ are observed at low temperatures and high pressures. The maximal deviation (∼2%) of the density of water saturated with CO₂ from the pure water was obtained at the borders of the tested ranges for both parameters, i.e. at 11 °C and 30 MPa.⁴¹

2.6 Viscosity of aqueous phase in biphase H₂O/CO₂ system

Similarly to the behavior of the density of water saturated with CO₂, the viscosity of this medium increases with the rise of pressure and mole fraction of the dissolved CO₂.⁴² At lower temperatures (around 23 °C as reported by McBride-Wright et al.⁴²) the maximum 10% increase of the viscosity of the mixture in comparison with the viscosity of pure water can be obtained. However, at higher temperatures this effect is only about 1% or slightly larger.

2.7 Dielectric constants of both phases in biphase H₂O/CO₂ system

One of the most important physical characteristics of a condensed matter is dielectric permittivity. This is the value that shows the relationship of the force of electrostatic interactions inside the medium as compared to the same in vacuum. When the dielectric constant is high the solvent is a polar one. When it is low the solvent is defined as a non-polar one.

The solvents with high dielectric constant promote the dissociation (for example of salts, or acids). The high dielectric constant facilitates the dissolution of ionic species. Substances that have high dielectric constant dissolve in the polar solvents (like water, with dielectric constant equal to 82 at 20 °C), while substances with low dielectric constant predominantly dissolve in non-polar solvents (like liquid or supercritical CO₂, with the dielectric constant around 1.5–1.6 at 20 °C and pressure in the range from 10 to 30 MPa).

Yet, water is miscible with some compounds having relatively low dielectric constants, e.g., tetrahydrofuran.^44,45 Liquid CO₂ behaves somewhat like a hydrocarbon solvent: it was shown by Hyatt⁴⁶ that small aliphatic and aromatic hydrocarbons, halocarbons, aldehydes, esters, ketones and low alcohols demonstrate good solubility in liquid CO₂. It is interesting, that CO₂ may efficiently extract water-soluble organics with low dielectric constants from their solutions in water in ternary mixtures.

It was proposed by Michels and Kleerekoper⁴⁷ that at high pressures the CO₂ molecules are “deformed”. This subsequently results in a change of the polarizability of the molecules. Indeed, it was shown⁴⁸ that high pressure can shift the electronic levels of the molecule and therefore all related physical constants depending on the electronic structure (including polarizability) will be affected too. The dielectric constant of the pure CO₂ increases with pressure (see Fig. 11) at fixed temperature.⁴⁹ Also one can notice from this figure that the temperature dependence is opposite: the dielectric constant of CO₂ decreases with temperature increase at the fixed pressure.

Fig. 11 Dielectric constant of pure CO₂ as a function of pressure. Data points were taken from ref. 49.

The well-known dependences of dielectric constant of pure water on the temperature⁵⁰ and pressure⁵¹ are shown in Fig. 12 and 13.

Fig. 12 Dielectric constant of pure H₂O as a function of temperature. Reprinted from ref. 50 copyright© 2003 with permission from Elsevier.

Fig. 13 Dielectric constant of pure H₂O as a function of pressure. Reprinted from ref. 51 in accordance with Creative Commons Attribution License.

Carbonic acid formed in the aqueous phase of the biphase H₂O/CO₂ system is an unusual acid with dielectric properties changeable by the external CO₂ pressure. Indeed, it is an aqueous (polar) medium, saturated with dissolved non-polar CO₂ molecules. The presence of the non-polar compound changes physicochemical properties of the solution (including dielectric constant). It is well known, that the dielectric constant of the mixture H₂O–non-polar gas is lower than the dielectric constant of pure water at constant temperature and pressure.^52,53

In the literature there is no direct experimental data about the dielectric constant of neither H₂O nor CO₂ phases of the biphase H₂O/CO₂ system. Nevertheless, the simplest estimations of the dielectric constants of the both phases can be carried out based on the assumption on the additive contributions (as a first approximation) in the polarizability of polar (H₂O) and non-polar (CO₂) components, as far as the interactions between dissimilar molecules (H₂O ↔ CO₂) in the system can be neglected. Therefore we can use Kirkwood approach (extension of the modified Onsager theory)^54,55 in order to estimate the dielectric constant of H₂O + CO₂ mixture, ε, for both CO₂ and H₂O phases:

where ρ is the density, P_i are the molar polarizations, x_i are the molar fractions and M_i are the molecular weights of the mixture components, i.e. of CO₂ and H₂O, respectively.

The molar polarizations of H₂O and CO₂ at given temperatures and pressures were extracted from the following equations:

The values of dielectric constant of pure water at different temperatures and pressures were taken from ref. 50, the values of dielectric constant of pure CO₂ at different temperatures and pressures were taken from ref. 49. The values of densities of pure water and pure CO₂ at different temperatures and pressures were calculated using “NIST Thermophysical Properties of Pure Fluids” software. The values of densities in the biphase H₂O/CO₂ system at different temperatures and pressures were taken from ref. 41. Eventually, the values of molar fractions of water dissolved in CO₂ and vice versa at different temperatures and pressures were taken from ref. 29.

As a result of the calculations we plotted the dielectric constant of water and CO₂ phases of the biphase H₂O/CO₂ system as functions of the pressure at three different temperatures (20 °C, 35 °C, and 50 °C), see Fig. 14 and 15. As one can see from the graphs in Fig. 14, the dielectric constant of water saturated with CO₂ is lower than that of pure water and it decreases with pressure. Whereas the dielectric constant of CO₂ phase saturated with water is higher than that of pure CO₂ and increases with pressure (see Fig. 15).

Fig. 14 Dielectric constant of the water phase of biphase H₂O/CO₂ system as a function of pressure.

Fig. 15 Dielectric constant of the CO₂ phase of biphase H₂O/CO₂ system as a function of pressure.

As a result, the water phase becomes less polar, when it is saturated with CO₂ and therefore the improved solubility of some organic species in water phase is expected, especially at high pressure. Contrariwise, the CO₂ phase in contact with water under pressure becomes somewhat more polar, and the solubility of more polar species in the CO₂ phase is also expected.

3. Organic synthesis in biphase H₂O/CO₂ system, carbonic acid

Peculiar physical properties of the biphase H₂O/CO₂ system, which were discussed above, eventually determine the course of the reactions to be performed in such media. The formation of carbonic acid in it allows one to perform acid-catalyzed reactions. The sensitivity of the main physical properties to the changes of pressure and temperature potentially ensures possibility of tuning or tailoring the reaction conditions during organic synthesis. In every practical case most optimal reaction conditions are still to be found (temperature, pressure, time of exposure). Yet, large amount of reports on different synthetic schemes has already been accumulated, thus certain analytics and generalizations are to be done and some general recommendations are to be formulated.

Indeed, the possibility to affect the composition, acidity, density, viscosity, dielectric constant of one or both phases in the biphase H₂O/CO₂ system by means of temperature and pressure variation was discussed above. The question is what would be general a priory recommendations to select most optimal pressure/temperature regime in a certain particular synthetic case from the viewpoint of obtaining better yields and/or selectivity. May such recommendations be even formulated? Unfortunately, at the present moment, some direct and straightforward answer is indeed still hardly to be formulated. Yet, certain knowledge of the main tendencies affecting general reactions courses is acquired.

When looking for proper temperature regime, one should take into account a competition of several driving forces. Firstly, the reactions are naturally activated with temperature due to Arrhenius mechanism. Secondly, solubility of CO₂ in water phase decreases, when temperature is increased starting since normal room values up to 60 °C and higher. Moreover, ability of the eventually formed H₂CO₃ molecules to dissociate is gradually suppressed with the temperature rise, see discussion above. The resulting reduction in acidity at elevated temperatures is unfavorable for acid-catalyzed reactions. This competition of (i) thermal activation and (ii) reduction in acidity may lead to an existence of certain temperature regions, which are either optimal or, on the contrary, non-optimal for a certain reaction. We will see below, that some acid-catalyzed reactions proceed particularly ineffectively when temperature is localized in a “dead valley” of around 60 °C. These reactions would proceed better, if the temperature is shifted to the higher or lower values. Yet, the exact localization of this unfavorable region may depend on the CO₂-saturating pressure as well. The situation is further complicated by the fact, that possible intermediates or byproducts may also have certain temperature regions of stability/instability.

When looking for proper saturating CO₂ pressure, one should take into account competition of some other factors. Of course, the higher the saturating pressure of CO₂ phase, the higher the acidity of the H₂O phase, equilibrated with it. Higher acidity is favorable for acid-catalyzed reactions. But it is very well known from the extraction experiments, that the addition of pressurized dense CO₂ may reduce the mutual solubility of water and organic solutes dissolved in it.⁵⁶ That means, in the biphase H₂O/CO₂ system too high CO₂-saturating pressures may results in effective extraction of products, intermediates and/or reagents from the water-rich phase to the dense CO₂/organic-rich phase. What does it possible mean in practice? The better extraction of products seems to be favorable as far as it facilitates separation. Efficient extraction of certain intermediates may also be favorable as far as it may introduce peculiar and in some cases unique (therefore, advantageous) selectivity to the reaction course. Yet, efficacy of extraction of reagents – before the reaction even occurs – means that excessive CO₂-saturating pressure may prevent the reagents from the access to the water phase and thus hinder the reaction to be catalyzed in it. The similar problem may appear for catalysts (if applied) with organic ligands, namely, if they are to be localized in a water phase, but yet may be extracted to a CO₂/organic phase if too high CO₂ pressures are applied. The catalysis in a non-polar CO₂/organic phase is typically somewhat suppressed. The efficiency of the extraction may be described by the corresponding partition coefficients for the two-phase system, which are to be determined experimentally in every particular case. Typically, the partition coefficient for an organic compound, soluble both in water and CO₂, increases manifold in favor of the CO₂/organic-rich phase with the CO₂ pressure increase. This is the reason why pressurized CO₂ is so efficient in extraction in general. That means, the variation of pressure may indeed become powerful tool to tailor the reaction course with respect to better yields and selectivity, but yet it is to be done with understanding of the ambivalent tendencies described here and the possibility to achieve two-fold outcome. On the other hand, efficiency in extraction of catalysts may provide a tool for their convenient separation and recovery for subsequent usage.^57,58

It is important to mention, that additional benefits of the biphase H₂O/CO₂ system as a reaction medium with respect to selectivity achievable may be brought on by CO₂-introdued chemical protection of certain reaction cites, such as, for example, aminogroups.

There are many examples in the literature, as we will see below, where the biphase H₂O/CO₂ system is used not due to the possibility of acid-driven catalysis, but just as a convenient biphase (polar/non-polar) system without a cross-contamination problem. Usage of some other catalysts besides carbonic acids may serve as a typical marker of such systems. The distribution of reagents, products and catalysts between the two phases may be different, which affects easiness and convenience of product separation, catalysts recovery and reagent replenishment. Indeed, in some cases possibility to perform continuous synthesis was demonstrated.

In the following subsections and sections we will describe and discuss in detail the available literature on organic synthesis reactions to be performed in biphase H₂O/CO₂ system. Besides, we will make a comparison with some other biphase systems, quite similar from the viewpoint of their chemical and physical properties.

3.1 Cyclization reactions

Cheng et al. studied cyclization reaction of citronellal to p-menthane-3,8-diol in biphase H₂O/CO₂ system at high pressures.⁵⁹ The product, p-menthane-3,8-diols (having several isomers), possess specific menthol-like smell and is widely used as a main active component of insects repellents. It occurs naturally being contained in rather small quantities in the leaves of Eucalyptus citriodora. Therefore, synthetic pathways of obtaining this valuable product from synthetic citronellal are of interest for the industry. Acid-catalyzed cyclization of citronellal is a typical way to produce p-menthane-3,8-diols. The synthesis has different stages, see Fig. 16. The first stage is isomerization of citronellal to isopulegol, which is a cyclic intermediate. The isopulegol reacts with water to form p-menthane-3,8-diols, though this reaction is reversible. The product may react further with citronellal in a condensation reaction thus forming p-menthane-3,8-diol citronellal acetals. Formation of acetal byproduct is not desirable if one is focused on obtaining p-menthane-3,8-diol, therefore selectivity of the synthetic process is important.

Fig. 16 Cyclization of citronellal to p-menthane-3,8-diol.⁵⁹

The authors found out that both the degree of conversion and the selectivity achieved are quite sensitive to pressure and temperature of the reaction. The conversion of the initial reagent, citronellal, is increasing (at the fixed temperature T: 120 °C, t = 2 h) with increase of the pressure of CO₂ saturating H₂O up to the value of ca. 7 MPa. But at further increase of the pressure the conversion started to decrease. This threshold value corresponds to critical pressure for CO₂ above which, according to phase diagrams, at the elevated temperatures sc CO₂ behaves as a supercritical liquid in distinct from a less dense supercritical gas with poorer dissolving power and thus extraction ability. Therefore, if this threshold value is passed, the efficient solubility of initial reagent, citronellal, in sc CO₂ phase is observed, where the reaction apparently does not occur, thus affecting conversion degree. Therefore, excessive CO₂ pressure prevents reagent access to the water phase. In other words, too highly pressurized CO₂ efficiently extracts the reagent from the reaction medium into the non-polar phase before the acid-catalyzed reaction occurs. Concerning selectivity with respect to the intermediate, isopulegol, the outcome of the p-menthane-3,8-diol was increasing (at the fixed temperature T: 120 °C, t = 2 h) up to the threshold value of 7 MPa. This increase may be explained by increase of acidity of the H₂O phase, which is essential for acidic catalysis of the reaction, with pressure increase of saturating CO₂. The subsequent decrease of the selectivity at even higher pressures may be explained by stronger affinity of isopulegol to sc CO₂ as compared to p-menthane-3,8-diols. Indeed, better extraction of the former from the H₂O to the CO₂ phase at these conditions may already be expected. Concerning the dependence of conversion and selectivity on temperature at the fixed pressure (1 MPa, t = 2 h), it was demonstrated that conversion was improved at the temperatures increase from 90 °C up to 110 °C. But it did not increase further in the temperature range of 110–130 °C. Apparently, this behavior is explained by the concurrence of two factors while temperature is increasing: (i) the reaction rate increase (activation) with (ii) the decrease of acidity of the H₂O phase. The selectivity to p-menthane-3,8-diols was deteriorated with temperature increase due to occurring of the reverse dehydration reaction of p-menthane-3,8-diols into isopulegols. For comparison, the authors studied catalytic behavior (at the same rather low temperature of 60 °C) of conventional acids, including sulfuric acid and acetic acid, which were superior to carbonic acid with respect to conversion and selectivity in the same reaction. The reaction proceeds markedly faster when sulfuric acid is applied as a catalyst. Yet, usage of carbonic acid simplifies the technological process (no neutralization or salt disposal to be done) whereas the reaction rate may be increased by means of temperature rise, if necessary.

Sun et al.⁶⁰ demonstrated how CO₂ may play an active role in reactions occurring in biphase H₂O/CO₂ systems. They performed typical reaction of cycloaddition of CO₂ to terminal epoxides (see Fig. 17), which is known to be catalyzed by a variety of cations and complexes.⁶¹ They compared many different Lewis bases as catalysts, almost all of which demonstrated improved activity in the presence of water (i.e. acidic medium when saturated with CO₂) in the cycloaddition reaction performed at 125 °C and saturating CO₂ pressure of 2 MPa, t = 1 h. The conversion, yield and reaction rate were also improved in the presence of water (i.e. carbonic acid). The authors proposed a concept that acidic water and the Lewis base play a synergic epoxy ring-opening effect as far as they coordinately attack the different parts of the epoxide, therefore presence of both is favorable for the rate of reaction. In general, moderate to excellent yields of products were achieved with high selectivity.

Fig. 17 Synthesis of cyclic carbonate from CO₂ and epoxide.⁶⁰

Lan et al.⁶² similarly studied cycloaddition of CO₂ to propylene oxide thus obtaining propylene carbonate, but they used graphene oxide as a metal-free, water-tolerant and highly efficient catalyst. The 96% yield and 100% selectivity were achieved at relatively mild conditions (CO₂ 2.25 MPa, 100 °C, 1 h). They emphasized particularly that the presence of a proper amount of water (i.e. carbonic acid) was essential for remarkably enhanced conversion of epoxide as compared with dry CO₂. Zinc-containing water-tolerant catalyst was also proposed for that reaction.⁶³ Ma et al.^64,65 reported simultaneous transformation of CO₂ together with glycerol to value-added products (glycerol carbonate, propylene glycol and propylene carbonate) via reaction with propylene oxide catalyzed by alkali metal halides (CO₂ 1–10 MPa, 105–125 °C, 1–2 h). Carbonation of epoxidized soybean oil even at atmospheric pressure of saturating CO₂ flow (100–140 °C) was also described.⁶⁶

Eghbali and Li⁶⁷ explored more benign approach to obtain the cyclic carbonates from CO₂ without using epoxides. They studied direct conversion of carbon dioxide and olefins into cyclic carbonates using organic base, 1,8-diazabicyclo[5.4.0]undec-7-ene, as a “CO₂ activator” in a presence of N-bromosuccinimide thus achieving metal-free synthesis. Applying 1.7–2.4 MPa of saturating CO₂ pressure they achieved almost complete conversion of different terminal alkenes and high selectivity for the desired product, cyclic alkenecarbonate, at the temperature of 60 °C (t = 3–6 h), see Fig. 18. The organic base serves as a deprotonating neutralizing agent. N-Bromosuccinimide may participate in the reaction forming bromohydrin but yet the final outcome of this byproduct is small. Therefore, when a bromohydrin intermediate is formed it apparently reacts further with CO₂. One should pay attention that the amount of CO₂ in this work and in some of the previous ones^60,62,66 is relatively small (saturating pressures are low), therefore CO₂ exists as a supercritical gas with poor dissolving power, therefore, no reaction should proceed in the CO₂ phase and no extraction by CO₂ should take place either. The same is true for several works cited below, where low pressures and correspondingly gas-like densities of CO₂ phase were explored.

Fig. 18 Conversion of terminal alkenes to cyclic alkenecarbonate in direct reaction with CO₂.⁶⁷

Gabriele et al.^68–70 focused on formation of maleic anhydrides and furan-2(5H)-ones from CO, CO₂, H₂O, and alk-1-ynes, see Fig. 19. It was found out that presence of both CO₂ (4 MPa) and small amounts of water were essential for the conversion. The reaction was performed in dioxane, at 80 °C, t = 24–64 h, in the presence of admixed gaseous CO (1 MPa), using PdI₂ and KI as catalyst. Both BuC≡CH and PhC≡CH were successfully converted to the mixture of corresponding anhydrides and furanones, see Fig. 19. For other examples on the reactions in the presence of additional admixed CO see the subsection below.

Fig. 19 Formation of maleic anhydrides and furan-2(5H)-ones from CO, CO₂, H₂O, and alk-1-ynes.⁶⁸

Li & Xie^71,72 performed Pd-catalyzed cyclotrimerization of alkynes in biphase H₂O/CO₂ system. Both aryl- and alkylacetylenes were converted into corresponding cyclotrimerization products regioselectively in high yields (CO₂ 1 MPa, room temperature, 24 h). Thus, environmentally benign protocol to synthesize some benzene derivatives was proposed. Similarly, tetra-ethyl pyrone (highly active pharmaceutical precursor) from hex-3-yne in water-in-CO₂ emulsions was produced.⁷³ Here nickel Triton X-100 surfactant complex was used as both an emulsifier and a water-soluble catalyst. The optimum yield of 69% was achieved at CO₂ pressure of 21 MPa (70 °C, 72 h). The authors mentioned easiness of the product separation as compared to other systems without any necessity for further purification. Timko et al.⁷⁴ studied Diels–Alder cycloaddition reaction between cyclopentadiene and methyl vinyl ketone in biphase H₂O/CO₂ system (8 MPa, 30 °C, 1 h) agitated with ultrasound, which improved the rate and the selectivity of the reaction. Chemical fixation of CO₂ by unsaturated alcohols using tert-butyl hypoiodite was reported even at atmospheric pressure.⁷⁵ Ma, Lu et al.^76,77 demonstrated chemical fixation of CO₂ reacting with 2-aminobenzonitriles to quinazoline-2,4(1H,3H)-diones in biphase H₂O/CO₂ system using choline hydroxide as a catalyst (CO₂ 2 MPa, 90 °C, 24 h).

3.2 Hydroformylation reactions

Other gases may be admixed to the biphase H₂O/CO₂ system in order to react with a substrate. Though dissolved CO₂ in water phase may affect solubility of other gases in it, good saturation of only one phase of the biphase system should be in general sufficient for other gases to effectively participate in the reactions. If CO₂ itself is in a gas state, complete miscibility of it with other gases is ensured due to just main properties of gases. The same is valid for supercritical CO₂. Moreover, solubility of gases in liquid CO₂ is also typically rather good.

McCarthy et al.⁷⁸ studied hydroformylation reaction in biphase H₂O/CO₂ system in the presence of syngas (H₂ + CO), see Fig. 20. Hydroformylation of alkenes to aldehydes by syngas is an industrially important process, which anyway is performed at high pressures (up to 30 MPa). The alkenes as reagents are relatively abundant and cheap in production. The obtained as products aldehydes are widely used in many fields of chemical industry and even may be further hydrogenated to alcohols and also in the biphase H₂O/CO₂ system. For example, the Ruhrchemie/Rhône Poulence (RCH/RP) oxo process of propene hydroformylation is based on usage of a rhodium catalyst with a water soluble ligand (Kuntz Cornils catalyst). But the separation of the product and the catalyst seems to be simpler if a CO₂-philic catalyst is applied in the biphase H₂O/CO₂ system. The authors prepared modified catalysts by mixing the rhodium precursor [Rh(CO)₂(acac)] and the CO₂-philic phosphine ligand 4-H²F⁶TPP, see Fig. 20. Using this catalyst they achieved complete conversion (100%) of the substrate to aldehyde at the following conditions of the reaction: the temperature of 60 °C, the partial CO₂ saturating pressure of 13 MPa, the syngas (CO & H₂) pressure of 2 MPa, t = 20 h. The possibility of efficient catalyst recycling was also demonstrated. The proposed reaction may be realized in a continuous process without repeated gas compression/decompression cycles as far as CO₂ with dissolved catalysts and reagents stays as the stationary phase whereas water layer accumulating the water-soluble product is disposed for the product removal (i.e., through a needle valve, as it was done by the authors).

Fig. 20 Hydroformylation of alkenes to aldehydes by syngas.⁷⁸

Schmid et al.⁷⁹ detected that addition of water (2–3%) is beneficial for ruthenium-catalyzed formylation of morpholine to N-formylmorpholine (H₂ 8.7 MPa, H₂ + CO₂ 21.5 MPa, 100 °C, 3 h). The authors believed that the addition of water suppressed solid carbamate formation thus enhancing the reaction rate.

Homogeneous hydroformylation of propylene was also performed in biphase H₂O/CO₂ system.⁸⁰ Rhodium phosphine complex with water-soluble ligand was used for catalysis. The temperature of 55 °C and pressure of 12–14 MPa (6 h) were reported as optimal conditions for high conversions to n-butyraldehyde and i-butyraldehyde (molar ratio between them of ca. 4.5). Possibility to reuse rhodium phosphine complex in successive synthetic cycles without significant losses of the catalyst was demonstrated. As compared to conventional hydroformylation, higher reaction rates and selectivity were detected.

Hallett et al.⁸¹ performed hydroformylation of 1-octene to aldehydes in the mixture of water with tetrahydrofuran (30 : 70 v/v). Tetrahydrofuran as a cosolvent was used in order to improve solubility of 1-octene in the reaction mixture. Rhodium complexes with ligands triphenylphosphine, triphenylphosphine monosulfonated sodium salt, and triphenylphosphine tris-sulfonated sodium salt were used as catalysts. Differently from what considered above, saturation of the reaction mixture with CO₂ was applied only after the reaction was completed in order to recover the catalysts and separate the product. Therefore, in this design CO₂ is used only as a “miscibility switch”⁸² (similar to a “solubility switch”⁸³) to facilitate efficient recovery and separation of the product and catalyst.⁸⁴ This important type of CO₂ usage is based on the fact that CO₂-saturation of the miscible binary mixture of tetrahydrofuran + water induces its phase separation and formation of a water-rich phase and a carbon dioxide + organic-rich phase, where the carbon dioxide amount increases with increasing pressure.^44,45,85 Therefore, the hydrophobic product tends to be extracted into this dense carbon dioxide + organic-rich phase, if sufficiently high pressure is applied (i.e., pressurized CO₂ is dense enough). The efficiency of the extraction is determined by corresponding partition coefficients, which increase with rising of CO₂-saturating pressure. On the contrary, the ligands of the catalysts tend to stay in the water-rich phase and the corresponding partition coefficients also markedly increase with rising of CO₂-saturating pressure. Eventually, possibility of convenient catalysts recycling was indeed demonstrated.⁸¹

3.3 Hydrogenation reactions

In this subsection we will mainly consider examples of hydrogenation of double bonds in biphase H₂O/CO₂ systems in the presence of admixed gaseous hydrogen. For reduction reactions in general see next subsection. Burgemeister et al.^86,87 particularly emphasized that usage of biphasic systems in homogeneous catalysis may be quite beneficial for catalysts recycling and product separation. Indeed, in their scheme there are two immiscible fluids, one contains dissolved catalyst and another contains products/substrates. They believe that such a system would be of a particular interest, where the catalyst resides in a non-polar phase and substrates/products are contained in a polar phase. Indeed, they have pointed out that many fine chemicals and biologically active compounds are highly polar and even water-soluble. Concerning the biphase H₂O/CO₂ systems, their scheme implies usage of CO₂-philic catalysts and water soluble products/substrates. Benefit of the biphase H₂O/CO₂ system as compared to other classical polar–nonpolar biphase systems is that there is no cross-contamination of the two phases, as it is often emphasized in relation to that.¹ In order to implement this scheme the authors used precursor complex [Rh(cod)₂][BF₄] (cod = 1,5-cyclooctadiene) and phosphine ligand tri-[3-(3,3,4,4,5,5,6,6,7,7,8,8,8-tridecafluoroctyl)phenyl]phosphane to prepare a catalyst soluble in sc CO₂. First example of successful implementation of their concept was rhodium-catalysed hydrogenation of itaconic acid to methylsuccinic acid, see Fig. 21. Itaconic acid is produced from carbohydrates in industrial quantities, whereas methylsuccinic acid may be useful in organic syntheses. It is well-known that itaconic acid can be converted to methylsuccinic acid, but proper hydrogenation catalyst is required, such as, e.g., platinum, which is expensive. Organometallic complex of rhodium is not too expensive; besides, it should demonstrate high partial activity and could be easily restored if the biphasic system synthesis is properly implemented.

Fig. 21 Rhodium catalyzed hydrogenation of itaconic acid.⁸⁶

The authors achieved complete conversion within several dozens of minutes at the temperature of 40 °C, CO₂ saturating pressure of about 10 MPa (taking into account the amount of CO₂, of water and the reactor volume), t = 1 h, and 3 MPa hydrogen partial pressure using magnetic stirrer to increase the interface surface. The authors made further step towards the illustration of importance of the biphase H₂O/CO₂ system for synthesis of biologically active compounds, which are known to be often chiral ones and moreover with different properties and behavior of their enantiomers. Therefore, enantioselective synthesis of enantiomerically enriched biologically active compounds is particularly interesting and the authors performed enantioselective hydrogenation in the biphase H₂O/CO₂ system using chiral phosphine–phosphite CO₂-philic catalyst. They detected that with this catalyst the hydrogenation of itaconic acid proceeds in an enantioselective manner. The same was demonstrated for hydrogenation of methyl-2-acetamidoacrylate as another benchmark substrate. Extremely high degrees of enantiomeric excess together with quantitative yield were detected for both model substrates studied. Thus, the possibility to synthesize almost enantiomerically pure products without any solvent impurities was successfully demonstrated.

Usage of emulsions may significantly increase interface area, which is favorable for reaction rates due to elimination of mass transfer limitations across the interface. Jacobson et al.⁸⁸ address the problem of catalyst separation and recovery, which may also be solved by the usage of emulsions. New opportunities are provided here also by unique properties of compressed and/or supercritical CO₂. Indeed, the pre-formed at high pressure emulsion can be broken by decreasing the pressure of CO₂, which results in formation of two separated phases with certain partition coefficients to be adjusted, thus facilitating eventually product separation and catalyst recycle. On the example of hydrogenation of alkenes catalyzed by water-soluble rhodium-phosphine complexes (RhCl(tppds)₃, tppds = tris(3,5-disulfonatophenyl)phosphine) they demonstrated drastic improvements in reaction rates when surfactants are applied to form the water-in-CO₂ emulsion. Different surfactants were successfully employed to stabilize the emulsion: (i) anionic perfluoropolyether ammonium carboxylate (PFPE COO–NH₄⁺, M_w = 740 and 2500 g mol⁻¹), (ii) cationic Lodyne 106A (C₆F₁₃(CH₂)₂SCH₂CH(OH)CH₂N⁺(CH₃)₃Cl⁻, M_w = 531.5 g mol⁻¹), and (iii) nonionic poly(butylene oxide)-b-poly(ethylene oxide) (M_w = 860-b-660 g mol⁻¹). The hydrogenation of styrene to ethylbenzene with complete conversion was achieved with a comparatively high rate (for a heterogeneous catalysis) in the presence of the surfactants at the temperature of 40 °C, CO₂ pressure of 28 MPa. The hydrogenated product is soluble in CO₂ at high pressure, which enables recovery procedure as well as CO₂ recycling. More hydrophobic 1-alkenes (1-octene, 1-decene, and 1-eicosene) were also successfully hydrogenated though with smaller rates. Yet, taking into account their much smaller solubility in water and not so dramatic reduction in reaction rate there are reasons to believe that the reaction occurs at the interface rather than in water bulk.

Ohde et al.⁸⁹ reported examples of successful catalytic hydrogenation of arenes with Rh nanoparticles in a water-in-supercritical CO₂ microemulsion using bis(2-ethylhexyl) sulfosuccinate (AOT) as a surfactant and perfluoropolyether phosphate as cosurfactant to stabilize the microemulsion. Further, Ohde et al.^90,91 studied the catalytic hydrogenations with Rh and Pd nanoparticles in the same water-in-CO₂ microemulsion in detail. Chalcone, nitrobenzene, and triphenylethylene were quantitatively hydrogenated to 1,3-diphenylpropane-1-ol, aniline, and 1,1,2-triphenylethane, respectively, at the temperature of 50 °C, the CO₂ + H₂ pressure of 20 MPa (partial hydrogen pressure of 1 MPa), t = 20–120 min, using 1-octanol as a cosolvent, see Fig. 22. No hydrogenation of the benzene ring of these organic compounds took place at these conditions. The hydrogenated products dissolved well in supercritical CO₂ and thus could be separated from the reaction mixture. The possibility to recover catalytic nanoparticles for subsequent catalytic cycles by pressure variation was demonstrated using the effect of phase separation of 1-octanol and sc CO₂, which are mixable only if the CO₂ pressure is above ca. 15 MPa.

Fig. 22 Catalytic hydrogenations of organic compounds with Pd⁰ nanoparticles stabilized in water-in-CO₂ microemulsion.⁸⁹

Liu et al.⁹² demonstrated that, on the contrary, CO₂-in-water emulsions may be sometimes beneficial as compared to water-in-CO₂ ones. They performed selective hydrogenation of citral catalyzed with palladium nanoparticles in the CO₂-in-water emulsion, see Fig. 23 (previously,⁹³ they studied the same process using complexes of transition metals with a phosphine ligand, triphenylphosphine, but in dry sc CO₂, i.e. without the presence of water). The emulsion was formed by using a nonionic surfactant of poly(ethylene oxide)–poly(propylene oxide)–poly(ethylene oxide) at the temperature of 50 °C, the CO₂ saturating pressure of 10 MPa and the hydrogen pressure of 4 MPa, t = 45 min. The conversion at certain conditions achieved values of 95% with high selectivity to citronellal (>75%). The further hydrogenated product, dihydrocitronellal, made 15%, besides traces of geraniol, nerol, isopulegol, as well as 3,7-dimethyl-1-octanol were also detected.

Fig. 23 Reaction pathways of citral hydrogenation.⁹²

When authors fixed the pressure (H₂ 4 MPa, CO₂ 10 MPa, t = 45 min) and varied the temperature in the range of 35–95 °C the conversion increased monotonically, whereas selectivity towards citronellal was decreasing, so more and more dihydrocitronellal was produced. This may be explained just by increase of rates of both hydrogenation reactions. When authors fixed the temperature (50 °C), t = 45 min, and varied the CO₂ pressure in the range of 6–14 MPa the conversion demonstrated non-monotonic volcano-type trend, whereas selectivity was more or less independent of pressure. The volcano-type behavior of the conversion may be explained by improving solubility of reagent, citral, in CO₂ phase (inside the micelles) at the initial pressure increase, but subsequent loosing of the micelles stability at excessively high pressures during further pressure increase. The authors performed comparison of their hydrogenation in CO₂-in-water emulsions with some other systems, including reverse water-in-CO₂ emulsions. The hydrogenation in CO₂-in-water emulsions markedly differed by higher conversion and higher selectivity towards citronellal. Higher efficiency was explained by the fact that in their reaction design both the reactants, citral and H₂, were efficiently dissolved in both the inner CO₂ core and outer surfactant shell of the micelles, thus ensuring two-side access of reagents to the Pd nanoparticles (i.e., from both the inner and outer surfaces of the emulsion droplets). Higher selectivity towards citronellal in their reaction design is favorable as far as citronellal is valuable product to be used in organic synthesis, see above. This higher selectivity towards citronellal is explained by better solubility of citronellal in sc CO₂ and thus its effective extraction by this fluid, which prevents further hydrogenation. The authors demonstrated also that catalytic Pd nanoparticles in CO₂-in-water emulsion could be recycled several times and still retain the same selectivity, though conversion is degrading with every cycle due to coalescence of the nanoparticles.

Further examples on hydrogenation reaction may be found in the next subsection.

3.4 Reduction reactions

Several relevant examples of reduction reactions were already mentioned in the previous subsection. In this subsection we will consider some other examples. In the literature on biphase H₂O/CO₂ system, hydrogen remains to be the most frequently used reducing reagent, apparently due to its ability to be removed completely from the reaction mixture after decompression, similarly to carbon dioxide. This is an important advantage as compared to some other reducing agents. Molecular or atomic hydrogen (the latter one to be generated in situ) may serve as a reagent for the reaction. The latter one is obviously much more active, but yet requires in situ generation, which produces species (salts) to be not so easily removed as compared to molecular hydrogen, CO₂ or carbonic acid.

Reduction of imines to amines formally is rather similar to hydrogenation of double bonds considered in the previous subsection. Ma et al.⁹⁴ studied reduction of imines to amines in biphase H₂O/CO₂ system in the presence of iron powder and zinc oxide, see Fig. 24. The reduction of imines and their derivatives to corresponding amines is a useful chemical transformation since the latter include many active ingredients or valuable building blocks of importance in pharmaceuticals and agrochemicals. Firstly, reduction of N-benzylideneaniline was tested as a model reaction. Combinations of Fe powder with different metal oxides were tested from the viewpoint of reaction promotion, among which combinations of Fe powder with ZnO, CuO or Fe₃O₄ gave certain yields of the product. The combination of Fe with ZnO was used in further experiments by the authors. When the reaction was performed at the fixed temperature of 40 °C, t = 15 h, and the varied pressure of CO₂ saturating water, non-monotonic dependence of both conversion and yield on the CO₂ pressure was observed. Indeed, both the conversion and yield initially increased when the saturating CO₂ pressure was increased from 2 up to 6 MPa (almost complete reaction) but further decreased at subsequent rise of the pressure from 6 up to 10 MPa. The authors explained the sluggish reduction at excessively high pressures (above 6 MPa) by a dilution effect of dense CO₂. Most probably, the dense pressurized CO₂ just prevents the reagent access to the water phase, i.e., it efficiently extracts reagent from the catalytic phase even before the reaction occurs. Whereas the initial improvement of the reaction behavior with the pressure rise may be related to increasing acidity of the water phase during saturating CO₂ pressure increase from 2 up to 6 MPa. Further, the authors proved that several other imines derived from different substituted aromatic aldehydes could be smoothly hydrogenated at high yields and selectivity.

Fig. 24 Reduction of imines to amines.⁹⁴

Jiang and Huang⁹⁵ studied synthesis of vicinal diamines via reduction of aromatic imines in CO₂ system with water presence, see Fig. 25. As a first model system they examined the reductive dimerization of N-benzylideneaniline at different temperatures and CO₂ pressures in the presence of water and Zn powder. Without CO₂ the reaction did not proceed, but already at the atmospheric CO₂ pressure (0.1 MPa) certain small degree of conversion was achieved. When the temperature was fixed at 80 °C, and the pressure was increased, the reaction yield was increasing until the pressure reached the value of about 8 MPa. At further increase of the pressure in the range of 8–29 MPa no influence on the reaction yield was detected. As optimal reaction time the authors selected the value of 6.5 h as far as for shorter times (3 h) somewhat smaller yields were detected. When the temperature was decreased from 80 °C down to 60 °C at the fixed pressure of 8 MPa the yield decreased even more noticeably.

Fig. 25 Synthesis of vicinal diamines via reduction of aliphatic and aromatic imines.⁹⁵

After the optimal reaction conditions were revealed (8 MPa, 6.5 h and 80 °C) the authors studied reductive couplings of a variety of other imines with pendant groups of different nature (both aliphatic and aromatic ones). They found out that the yields achievable are affected by electronic effects of the substituents in imines. When R¹ pendant group was more electron-rich whereas R³ pendant group was more electron-deficient, the yields were smaller, and vice versa.

Reduction of aldehydes to the corresponding alcohols formally also is similar to what is considered in the previous subsection. Bhanage et al.⁹⁶ performed hydrogenation in biphase H₂O/CO₂ of an α,β-unsaturated aldehyde, cinnamaldehyde, into corresponding alcohols with special focus to obtain unsaturated alcohol. They mentioned easy catalyst/product separation and catalyst recycling as the main benefits of the reaction system. Complexes of Ru, Rd, Pd with a water-soluble P(C₆H₄SO₃Na)₃ (TPPTS) ligand were applied as catalysts. The reaction was performed at the temperature of 40 °C, the CO₂ saturating pressure of 140 MPa and the hydrogen partial pressure of 40 MPa, t = 2 h. Relatively poor conversion degrees (38–44%) were demonstrated, but high selectivity (96–99%) towards unsaturated alcohol, cinnamyl alcohol, was achieved with Ru-based catalysts.

Fujita et al.⁹⁷ also studied selective hydrogenation of trans-cinnamaldehyde in biphase H₂O/CO₂ system in the presence of hydrogen (CO₂ 12–20 MPa, H₂ 2 MPa, 70 °C, 2 h) using water-soluble ruthenium/phosphine complex catalysts. The selectivity for cinnamyl alcohol as well as total conversion decreased after the first run (i.e., after initial passivation of the complex with cinnamyl alcohol product) but on the subsequent runs remained more stable. Successful reduction of several aldehydes in biphase H₂O/CO₂ system was reported also in the presence of Fe (CO₂ 8 MPa, 150 °C, 24 h).⁹⁸ Hiyoshi et al.⁹⁹ demonstrated that addition of CO₂ and H₂O enhances hydrogenation of acetophenone to 1-phenylethanol over an activated carbon-supported palladium catalyst (CO₂ 0.8 MPa, H₂ 3 MPa, 40 °C, 45 min). Successful hydrogenation of 5-hydroxymethylfurfural to tetrahydro-5-methyl-2-furanmethanol (CO₂ < 10 MPa), 2,5-dimethylfuran (CO₂ 10 MPa) or 2,5-dimethyltetrahydrofuran (CO₂ > 10 MPa) was reported with very high selectivity in biphase H₂O/CO₂ system in the presence of hydrogen over a Pd/C catalyst (H₂ 0.2–1 MPa, 80 °C, 2 h).¹⁰⁰

Liu et al.¹⁰¹ developed one-pot production of industrially valuable diketone derivatives from biomass-produced carbohydrates by means of ring-opening Pd@C-catalyzed hydrogenation/hydrolysis in biphase H₂O/CO₂. Readily obtained from carbohydrates dimethylfuran was completely and selectively hydrolyzed to the corresponding 2,5-hexanedione (CO₂ 4 MPa, 150 °C, 15 h). Further, 2-methyl-5-hydroxymethylfuran or 2,5-dihydroxymethylfurane were converted into the mixture of 2,5-hexanedione and 2-methyl-5-hydroxymethyltetrahydrofurane, or into 1-hydroxyhexane-2,5-dione, respectively, in the presence of hydrogen (CO₂ 3 MPa, H₂ 1 MPa, 150 °C, 15 h). Moreover, 5-hydroxymethylfurfural was also converted to 1-hydroxyhexane-2,5-dione (CO₂ 3.5 MPa, H₂ 0.5 MPa, 120 °C, 15 h). Eventually, catalytic conversion of fructose and inulin to 1-hydroxyhexane-2,5-dione was also demonstrated. Thus, the approach for the production of valuable diketone derivatives from biomass products was illustrated.

It is known that reduction of aldehydes to alcohols efficiently proceeds in acidic aqueous media in the present of Zn powder. Li et al.¹⁰² revealed that in CO₂ system with water presence this reaction is characterized by good yields and excellent selectivity. Reduction of p-tolualdehyde to 4-methylbenzyl alcohol was initially tested as a first typical model reaction, see Fig. 26.

Fig. 26 Reduction of p-R-aldehydes to 4-R-benzyl alcohols.¹⁰²

Without CO₂ the reaction does not occur, yet, even at the atmospheric pressure of saturating CO₂ (0.1 MPa) some conversion is observed, see Fig. 26. At the temperature of 65 °C the reaction yield demonstrates non-monotonic volcano-type trend: initially increasing and achieving maximum at the pressure of 8 MPa it is decreasing at further pressure increase, particularly above 15 MPa (t = 6 h). This typical non-monotonic behavior should be again explained by effective extraction of organic reagent by the compressed dense CO₂ from the catalytic water phase before the reaction even occurs (that means suppression of the reagent access to water phase). The authors also compared efficiency of the reaction for different aldehydes (t = 6–24 h), including aliphatic and aromatic ones. It was found out that the aromatic aldehydes demonstrate significantly larger conversion degrees as compared to aliphatic ones.

Shirai et al.¹⁰³ also described catalytic hydrogenation of benzaldehyde to benzyl alcohol in biphase H₂O/CO₂ using charcoal-supported palladium catalyst. The reaction was performed at the temperature of 40 °C, at the hydrogen pressure of 1 MPa and the CO₂ saturating pressure of 3 MPa, t = 0.5 h. The conversion achieved was 53%, which is to be compared with 46% conversion in the absence of CO₂ (i.e. in pure water). The selectivity towards benzyl alcohol was 98.7%, the rest was accounted for toluene, the product of complete hydrogenation. When varying the saturating CO₂ pressure in the range of 0.1–15 MPa, non-monotonic dependence of yields on the carbon dioxide saturating pressure was revealed with maximum observable at ca. 5 MPa value of the CO₂ saturating pressure (at 3 MPa of the hydrogen partial pressure). As above, this typical non-monotonic behavior may be explained by effective dissolution (extraction from water) of the reactant, benzaldehyde, in sc CO₂ at the excessive pressures (prevention from the reaction).

Ma et al.¹⁰⁴ reported reduction of sulfoxides and pyridine-N-oxides in biphase H₂O/CO₂ system in the presence of metal powders (Fig. 27). The reduction here means deoxygenation of sulfoxides and pyridine-N-oxides to the corresponding sulfides and pyridines, which is a rather important synthetic route for organic synthesis and biological chemistry. Indeed sulfoxides may be used as chiral synthetic building blocks in asymmetrical synthesis to control the stereochemistry of the product in the key step of the total synthesis of biologically active compounds. But their chiral sulfinyl group is to be eliminated after stereoselective induction. Similarly, pyridine-N-oxide and other heterocyclic oxides may serve as a source of functionalized pyridine derivatives after the functionalization and subsequent reduction. Therefore, pathways of their reduction reaction in biphase H₂O/CO₂ system are worthwhile to be studied in detail.

Fig. 27 Reduction of sulfoxides and pyridine-N-oxides.¹⁰⁴

Several metal powders were tested to promote the reaction, among which Fe and Zn provided sufficiently high yields of the reduction. With Fe powder the reaction proceeds fast with high yields in the presence of water at the temperature of 80–100 °C and the CO₂ saturating pressure in the range of 0.5–8 MPa, t = 10 h. No reaction is observed without saturating CO₂ (in pure water). Interesting to note, that high pressures were actually excessive: 0.5 MPa of CO₂ saturating pressure was already quite sufficient to achieve reasonable rate and yields. That means the reaction procedure and the equipment required could be essentially simplified. Moreover, the reduction was again proven to be sluggish at the higher pressure (8 MPa) apparently due to the same typical extraction effect of dense CO₂, as discussed above. Many different sulfoxides and pyridine-based N-oxides were successively reduced in the biphase H₂O/CO₂ system in the presence of Fe powder to the deoxygenated products in moderate to excellent yields. Concerning the reaction pathway, apparently atomic hydrogen being generated in carbonic acid in the presence of metallic Fe serves as an efficient reducing agent whereas the Fe species may serve as mediators of the hydrogenation.

Yoshida et al.¹⁰⁵ performed hydrogenation of benzonitrile in biphase H₂O/CO₂ in the presence of hydrogen with a Pd@Al₂O₃ catalyst. Benzylamine, dibenzylamine, and toluene were detected as possible products with almost complete and rather selective conversion to the first of them at optimized conditions (H₂ 2 MPa, CO₂ 7 MPa, 50 °C, 4 h) with no deactivation of the catalyst. The authors explained enhanced selectivity to the primary amine by its transfer into water phase, which prevented it from reacting with an intermediate of imine into the secondary amine. In addition, they believed that interactions of CO₂ molecules with the primary amine could also increase its selectivity. Selective hydrogenation of benzonitrile with Ni@Al₂O₃ was compared in several biphase systems including H₂O/CO₂ (H₂ 4 MPa, CO₂ 10 MPa, 80 °C, 1–12 h).¹⁰⁶ The decreased yield of benzylamine in the latter case was detected because the carbonic acid favors the elimination of NH₃. Further, Yoshida et al.¹⁰⁷ studied hydrogenation of phenol to cyclohexanone/cyclohexanol in biphase H₂O/CO₂ in the presence of hydrogen with a Pd@Al₂O₃ catalyst. They mentioned the problem of Pd leaching due to carbonic acid formation as well as catalyst poisoning by CO adsorption, which is formed from H₂ and CO₂. Hydrogenation of benzyl cyanide to 2-phenylethylamine in multiphase n-hexane + H₂O + CO₂ system was performed in the presence of hydrogen with a Pd@Al₂O₃ catalyst (H₂ 2 MPa, CO₂ 3 MPa, 50 °C, 1 h).¹⁰⁸ The authors noticed synergistic effects when both H₂O and CO₂ are included in the reaction medium. Eventually, Yoshida et al. summarized features of selective hydrogenation reactions with supported catalysts for different substrates in several biphase systems including H₂O/CO₂.¹⁰⁹

Reduction of nitroarenes to anilines (aromatic amines) is a well-studied and widely used process (Zinin reaction is a typical example to be mentioned here). Aromatic amines are extensively applied as building blocks in the synthesis of pharmaceutical products, dyestuffs and polymers. Jiang and Dong¹¹⁰ studied complete reduction of nitroarenes to anilines in CO₂ in the presence of water and Zn powder. Hydrogenation of nitrobenzene to aniline was used as a first model system. No reaction occurred in the absence of CO₂, i.e. in pure water with Zn powder. Already atmospheric pressure of CO₂ saturating H₂O (by bubbling) was enough to achieve yield of ca. 24% at 80 °C, t = 12 h. Increasing the CO₂ pressure at the fixed temperature (80 °C, t = 12 h) resulted in initially increased yield up to 97% at 8 MPa due to increasing acidity. But at even higher pressures the yield achievable was smaller: 86% at 12 MPa and 80% at 15 MPa. Apparently, the excessive CO₂ pressures result in dissolution of the reagent in this phase and its extraction from water thus preventing its contact with the catalytic acidic phase in the presence of atomic hydrogen being generated in H₂O phase from Zn powder. When the temperature was reduced at the fixed pressure of 8 MPa, smaller yields were obtained with smaller temperatures (77% at 60 °C and 65% at 45 °C, t = 12 h) indicating on activation mechanism of the reaction. Reductions of other nitrobenzenes with different pendant groups, –CH₃, –Cl, –Br, –F, –I, –NH₂, –OH, –COOH, were further investigated. Some of these functional groups may also be reduced, but the authors detected that at these conditions (8 MPa, 80 °C, 12 h) the reaction was mainly chemoselective (with an exception for the –I group) and only the nitro-group was affected. The yields were different for different pendant groups but mainly above 90%.

Besides the complete reduction of nitroarenes to anilines considered above, intermediate reduction of nitroarenes to N-aryl hydroxylamines is also important for the synthesis of fine chemicals and pharmaceuticals. Liu et al.¹¹¹ performed this reaction in the biphase H₂O/CO₂ system, i.e. in carbonic acid in the presence of Zn powder (Fig. 28). Reduction of nitrobenzene to N-phenylhydroxylamine was tested first as a model reaction. Almost no reduction took place without CO₂ saturating H₂O but already atmospheric CO₂ pressure (0.1 MPa) was enough to accelerate the reaction dramatically. If the temperature was fixed at 40 °C, t = 3 h, the pressure increase from 0.1 to 0.5 MPa resulted in the increase of both conversion (complete) and selectivity towards N-phenylhydroxylamine (76%). But at further pressure increase from 0.5 MPa up to 5 MPa, t = 3 h, the conversion remained almost complete, whereas selectivity towards N-phenylhydroxylamine gradually decreased down to 14% with corresponding increase of the yield of aniline, the product of a complete reduction. The same trend of increased production of aniline instead of N-phenylhydroxylamine was observed with increase of amount of Zn powder or prolonged times of the reaction. When the CO₂ saturating pressure was fixed (P = 0.1 MPa), t = 1 h, and the temperature varied, the dependence of both conversion and selectivity towards N-phenylhydroxylamine demonstrated non-monotonic dependence on temperature similar to each other. The increase was observed in the temperature range of 0–25 °C, then decrease in the range of 25–55 °C, then increase again in the range of 60–75 °C with evident saturation (no noticeable changes) in the range of 80–100 °C. This complex behavior may be explained by a concurrence of at least two different trends: (i) activation of reaction with temperature, (ii) non-linear decrease of acidity of carbonic acid with temperature in the temperature range of interest. Finally, the authors demonstrated that in the biphase H₂O/CO₂ system it is possible to reduce selectively other nitroarenes, i.e. ones substituted with other reducible functional groups, such as –COCH₃ and –CN (Fig. 28). Later on Liu et al.¹¹² repeated the same reactions in the same media at additional influence of ultrasound, which improved yield, shorted reaction times, and decreased consumption of Zn powder.

Fig. 28 Reduction of various substituted nitrobenzenes.¹¹¹

Gao et al.¹¹³ studied reduction of nitroarenes with pendant groups to corresponding aromatic amines. The goal here was to develop a reaction process in such a way as to reduce only the nitro-group selectively and thus to preserve other potentially reducible groups intact, such as chloro-, carbonyl- and cyano-groups. Indeed, aromatic amines carrying such groups serve as important building blocks in the synthesis of many fine chemicals, including pharmaceutical and agricultural ones. Completely reduced products could be easily obtained from other more simple substrates. As a first example they studied reduction of p-chloronitrobenzene by iron powder in the biphase H₂O/CO₂ system. Indeed, after the reduction only p-chloroaniline was detected and not any traces of aniline, possible product of dechlorination. That means that the chloro-group was indeed preserved during the reduction process. Yet, almost no reaction was observed without saturating CO₂, i.e. in pure water with iron powder. When the reaction temperature was fixed at 120 °C (t = 8 h) and the pressure was varied, they detected initial monotonic increase of the p-chloronitrobenzene yield with the pressure increase up to the range of 6–8 MPa, where almost complete (∼100%) yield was detected. At further pressure increase in the range of 8–16 MPa the yield was somewhat decreasing by 11–13%. This is typical again: at such excessive pressures the substrate is effectively extracted by pressurized sc CO₂ phase from water, which prevents its contact with iron-containing phase and conversion. Increased iron consumption due to acidity of the water phase at high pressures of saturating CO₂ may also contribute to the observed behavior. When the pressure was fixed at 7 MPa (t = 8 h) and the temperature was varied in the range of 8–150 °C they observed complex non-monotonic behavior. Initially, the product yield was increasing with temperature increase up to the first maximum (∼55%) at the temperatures around 30 °C but then demonstrated subsequent pronounced reduction with marked minimum (∼7%) at the temperatures of about 60 °C. At subsequent temperature increase the yield was growing again up to the second maximum (∼100%) at the temperatures around 120 °C but then again demonstrated subsequent certain reduction down to 75% at the temperatures about 150 °C. This complex behavior is explained mainly by competition of two major tendencies: (i) activation of the reaction rate with temperature rise; (ii) non-monotonic dependency of acidity of the water phase on temperature at fixed pressure. At the temperatures above 120 °C additional contributing factor to the decrease of p-chloronitrobenzene yield may be related to increased iron consumption. The authors also performed successful reduction of three other nitroarenes, which were converted to corresponding amines with preservation of pendant groups thus proving high chemoselectivity of this reaction system.

Cheng et al.¹¹⁴ comparatively studied hydrogenation of o-chloronitrobenzene to o-chloroaniline in biphase H₂O/CO₂ system over Pt@C and Pd@C as catalysts. They proved that combination of H₂O and CO₂ as a reaction medium is more efficient than pure H₂O, dry CO₂ or H₂O/n-heptane biphase system. The authors detected enhancement of reaction rate and the maximum conversion at saturating CO₂ pressure of 9 MPa (H₂ 4 MPa, 35 °C and at even rather short reaction times of 15–50 min). High selectivity was also achieved though detected intermediates included o-chloronitrosobenzene, dichloroazoxybenzene, dichloroazobenzene, and dichlorohydrazobenzene.

3.5 Coupling reactions

Li et al.^115–117 implemented reductive Ullmann coupling of various aromatic halides in biphase H₂O/CO₂ system, see Fig. 29. As a first model system they used coupling of iodobenzene to biphenyl activated by Pd@C and Zn powder at room temperatures, t = 8 h. It was found out that without saturating CO₂ the reaction was slow and inefficient in water. Yet, already saturation of water with CO₂ at atmospheric pressure (0.1 MPa) by means of simple bubbling resulted in increase of the yield of the biphenyl up to 41%. Further increase of the pressure of saturating CO₂ up to 1 MPa resulted in complete conversion (100%) of iodobenzene to biphenyl. It is interesting that again at yet even higher pressures, e.g. of 6 MPa, typical decreased yield (only 65%) of biphenyl was obtained with the remaining portion (34%) of benzene, the product of iodobenzene complete reduction. Nevertheless, the possibility to obtain at properly chosen parameters the yields of biphenyl close to 100% indicates on unique and peculiar selectivity of the reaction in biphase H₂O/CO₂ system.

Fig. 29 Palladium-catalyzed reductive Ullmann coupling of aromatic halides.¹¹⁵

To demonstrate general nature of the approach the authors successfully performed the same coupling of other various aromatic halides bearing different pendant and functional groups. The reaction occurrence depended on the structure of the halides, i.e. on both the nature and position of substituents on the aromatic ring. In general, symmetric biaryls were obtained with high yields (almost complete conversion, high selectivity) and reasonable rates from aromatic iodides and, to a bit lesser degree, from aromatic bromides. Aromatic chlorides were generally less reactive, yet, increase of the reaction times up to 36 h and in some cases up to 96 h allowed authors to achieved comparably high yields of the coupling as well. Later on Zhang et al.¹¹⁸ proposed Pd/Ph on mesoporous silica nanoparticles SBA-15 as a catalyst for this type of Ullmann reaction.

Heck coupling reaction between a halide with an alkene is rather important for organic synthesis due to its possible stereoselectivity, etc. Bhanage et al.¹¹⁹ performed Heck coupling of iodobenzene in biphase H₂O/CO₂ system with butyl acrylate and with styrene. They used water soluble triphenylphosphine trisulfonate sodium salt (TPPTS) as a ligand for Pd-based catalyst. The reaction rates were very low without water, i.e. in pure sc CO₂. With addition of water somewhat higher, but yet relatively poor, conversions in the range of 8–18% were achieved at the temperature of 60 °C for t = 17 h. Products, an organic phase, were easily to be separated from the catalyst soluble in water phase, thus facilitating its recovery and subsequent usage. It is interesting, that conversions were higher if ethylene glycol was used as a co-solvent instead of water thus serving as a gas-expanded reaction medium (see below). Eventually, Bhanage et al. in ref. 120 reviewed and summarized their activity on multiphase Heck reactions using various catalysts, including the ones for biphasic systems.

Ye et al.¹²¹ performed Pd-catalyzed Suzuki coupling reaction in water-in-CO₂ microemulsions, stabilized by nonionic surfactants of TMN series. For example, iodobenzene and 4-methylbenzeneboronic acid were cross-coupled in this system to produce p-methyl biphenyl (CO₂ 12 MPa, 80 °C, 25 h). Some other aryl halides (substituted ones) were also successively coupled with arylboronic acids at the same conditions. Comparing with other systems (pure water, water–cyclohexane microemulsion), the authors mentioned synergistic effects from the presence of both CO₂ and H₂O.

3.6 Rearrangement reactions

Liu et al.¹²² studied Bamberger rearrangement of N-phenylhydroxylamines to p-aminophenols in biphase H₂O/CO₂ system, see Fig. 30. This reaction is rather important as far as the reagents, N-phenylhydroxylamines, may easily be synthesized from nitrobenzenes, whereas the products, p-aminophenols, as well as some of their derivatives, serve as useful building blocks in industrially important chemical processes, including pharmaceutics and synthesis of dyes. When the temperature was fixed at 100 °C (t = 1 h) the variation in pressure was demonstrated to affect the reaction significantly. Without CO₂ no reaction occurred in the pure water at 100 °C. At CO₂ saturating pressure increase up to 2 MPa both the conversion of N-phenylhydroxylamine and yield of p-aminophenol were increased monotonically in a symbate manner, particularly fast in the range of up to 0.5 MPa. After exceeding 4 MPa of the CO₂ saturating pressure, no further increase in either conversion of N-phenylhydroxylamine or yield of p-aminophenol was observed, which both demonstrated saturation behavior at the values of 100% and 80%, respectively. There were also stable amounts of byproducts detected, such as aniline, azoxybenzene and nitrosobenzene, accounting for the remaining 20% (Fig. 30). This reaction behavior may be related to the acidity of carbonic acid solutions, which also demonstrates certain saturation at higher pressures of CO₂ applied.

Fig. 30 Bamberger rearrangement of N-arylhydroxylamine to p-aminophenol, main reaction, and concomitant side reactions.¹²²

When the CO₂ saturating pressure was fixed at 4 MPa value (t = 1 h), both conversion of N-phenylhydroxylamine and yield of p-aminophenol were increasing with temperature increase in the range of 40–100 °C, being particularly low at the lower temperature of the indicated range. But at the temperature of 100 °C and above the corresponding conversion and yield stabilized at the values of 100% and 80%, respectively. This behavior may indicate that certain activation of the rearrangement reaction is required.

The authors demonstrated also that N-arylhydroxylamines substituted with –CH₃ and –Cl groups may be efficiently converted to the corresponding p-aminophenols with even higher selectivity (i.e., smaller amounts of byproducts) due to electronic effects. Indeed, for electron-donating CH₃ pendant group the yield of 4-amino-3-methylphenol achieved as high value as 92%.

Zhang et al.¹²³ performed disproportionation reactions of aryl alcohols (1-phenylethanol, 1-(4-methoxyphenyl)ethanol, 1-indanol) in biphase CO₂/H₂O system. They proved that addition of CO₂ to water phase increases the reaction rate significantly (CO₂ 6 MPa, 100 °C).

3.7 Substitution/addition reactions

Tundo et al.¹²⁴ studied formation of diazonium salts from primary aromatic amines in biphase CO₂/H₂O system and their transformation to corresponding aryl iodides by potassium iodide, see Fig. 31. This is a radical-nucleophilic aromatic substitution reaction. Diazonium salts with aryl groups are important building blocks in the organic synthesis, particularly, e.g., of azo dyes. The widely adopted method for their obtaining is reaction of aromatic amines (anilines) with nitrous acid, which is to be generated in situ by reaction between sodium nitrite and some mineral acid. Therefore, it is only natural to try to replace the mineral acid with the self-neutralizing carbonic acid. When the authors performed the reaction at room temperature (∼25 °C), t = 6 h, employing p-anisidine as a substrate the yield of aryl iodides was 85–95%. The yield was particularly high (95%) when liquid CO₂ was formed in the reactor at the saturating CO₂ pressure of 6.5 MPa as compared to the somewhat smaller yield (85%) when gaseous CO₂ was explored at the saturating pressure of 3 MPa. That may be attributed to the higher density of liquid CO₂ and, as a consequence, to its higher dissolving power towards aryl iodides, the product. Formation of biphenyls was not detected, though they were formed if smaller reaction temperatures (∼0 °C) were chosen. Formation of phenols, possible by-product, was not ever observed. In distinct from classical scheme of Sandmeyer reaction no catalysts such as copper salts were required.

Fig. 31 Formation of diazonium salt and its reaction with I⁻.¹²⁴

As far as in this reaction system aryl iodides, the product, are well soluble in liquid CO₂ phase, whereas anilines, the initial substrate, are soluble in liquid CO₂ only slightly and diazonium salts, the intermediate, are insoluble in liquid CO₂ and soluble in water phase only, separation of the product from the reaction mixture containing the intermediate may be performed using simple venting procedure with CO₂. Further, several other primary aromatic amines (aniline, p-nitroaniline, p-chloroaniline) were tested in the reaction as well. In all cases the conversions toward aryl iodides were higher when CO₂ presented in the reactor as a liquid (6.5 MPa) and not as a gas (3 MPa). Yet, in general the yields were noticeably smaller as compared to the p-anisidine case. The authors tested possibility to convert also aliphatic amines using n-octylamine as an example, and though the conversion was complete, rather poor selectivity was observed. Therefore, this reaction scheme is very promising for synthesis of aryl halides but hardly suitable for preparation of alkyl halides.

Jacobson et al.^125,126 considered the nucleophilic substitution reaction between hydrophobic benzyl chloride and a hydrophilic nucleophile, KBr, in water-in-carbon dioxide (w/c) and carbon dioxide-in-water (c/w) microemulsions and emulsions (Fig. 32).

Fig. 32 Nucleophilic substitution reaction between benzyl chloride and KBr.¹²⁶

For stabilization of the emulsion¹²⁶ they used anionic perfluoropolyether ammonium carboxylate, nonionic poly(dimethylsiloxane)-g-poly(ethylene oxide) and poly(butylene oxide)-b-poly(ethylene oxide) surfactants. The reaction was performed at the CO₂ saturating pressure of 28 MPa, t = 5 h, at the temperature of 45 °C and 65 °C, for comparison. The achieved yields were very poor, in the range of 5–7%, at 45 °C, but significantly larger, in the range of 29–47%, at 65 °C. The typical benefit of the demonstrated approach is that the emulsion can be easily broken by decreasing the pressure to separate the water and CO₂ phases, thus facilitating product recovery and CO₂ recycle.

Blanchard & Brennecke mentioned that esterification of acetic acid with ethanol proceeded better in the presence of CO₂ (6 MPa, 60 °C).¹²⁷ Zhang et al.¹²⁸ developed the process of urea transamidation with amine to urea derivative in biphase H₂O/CO₂ system (5–6 MPa, 120–185 °C, 10 h) without any catalyst (besides the in situ generated carbonic acid). Cenci et al.¹²⁹ performed metal-mediated Barbier allylation in ultrasonically agitated biphase H₂O/CO₂ system. A variety of substituted benzaldehydes were converted by allyl halides to corresponding homoallylic alcohols in moderate to high yields in the presence of Zn (CO₂ 8–12 MPa, 30–60 °C, 2 h).

3.8 Halogenation reactions

Li et al.¹³⁰ studied diiodination of alkynes in biphase H₂O/CO₂ system, see Fig. 33. The products of the diiodination reaction, 1,2-diiodoalkenes, are useful intermediates in organic synthesis. Diiodination of phenylacetylene was used as a model system. The authors detected that without CO₂, i.e. in pure water, the reaction was slow: 19% yield of the product, trans-1,2-diiodophenylethylene, was obtained for 8 h (room temperature). In the presence of CO₂ (1 MPa) complete conversion with quantitative yield of trans-1,2-diiodophenylethylene was achieved for 3 h (room temperature). Further, diiodination of other alkynes (C₈H₁₇C≡CH, C₃H₇C≡CC₃H₇, CH₃C≡CCOOMe, HC≡CCOOEt) to corresponding trans-1,2-diiodoalkenes was carried out successfully with quantitative yields as well. Diiodination products were readily isolated in water phase after the release of CO₂.

Fig. 33 Diiodination of alkynes.¹³⁰

3.9 Hydrolysis reactions

Rayner¹³¹ mentioned possibility to perform carbonic acid-catalyzed hydrolysis of ketals, acetals, and epoxides at the pressure of CO₂ saturating water being of 2 MPa and the temperature in the range of 50–65 °C, t = 4–48 h. The mentioned CO₂ pressure is rather low, therefore CO₂ has low density and exists in a sc gas state. Possible benefits of the selection of the relatively low pressure may be related to the situations, when it is not desirable that organic compounds, reagent or products are to be dissolved in sc CO₂ and extracted from the water phase. We saw above that excessive CO₂ pressures may often reduce the reaction speed. Different ketals, acetals, and epoxides can be hydrolyzed completely at these conditions, see Fig. 34, whereas the reaction almost does not proceed at the same temperatures in the absence of saturating CO₂, i.e. at pH neutral conditions (pure water). This proves catalytic activity of carbonic acid in the hydrolysis reaction of ketals, acetals, and epoxides.

Fig. 34 Carbonic acid-catalysed hydrolysis of ketals, acetals, and epoxides.¹³¹

Timko et al.^132,133 studied hydrolysis of benzoyl halides in biphase H₂O/CO₂ system as a model reaction, see Fig. 35. Their focus was to investigate the reaction/mass transport dynamics.¹³² The reagents, benzoyl halides, are extremely poorly soluble in water, but they are highly soluble in CO₂ phase. Yet, the hydrolysis reaction should be suppressed in the CO₂ phase due to poor solubility of water in CO₂ and known low reaction rate of hydrolysis in nonpolar solvents, such as CO₂. Therefore, mass transfer of reagents from CO₂ phase towards water phase should affect significantly the reaction rate.

Fig. 35 Hydrolysis of benzoyl halides.¹³²

The reaction was performed in the presence of water at the CO₂ saturating pressure of 6.5 MPa and the temperature of 30 °C. Ultrasound was applied to homogenize biphase H₂O/CO₂ mixture in the form of emulsions (both of H₂O-in-CO₂ and CO₂-in-H₂O). The measured diameters of both water and carbon dioxide droplets in both types of emulsions were about 10 μm. Seven benzoyl halides were tested in the hydrolysis reactions. It was found out that ultrasonic agitation increases hydrolysis rates in all cases up to 200 times. According to theoretical model provided, the diffusion resistance in water phase effectively controls the reaction rate. Sonication results in increased reaction rate due to increase the carbon dioxide/water interfacial area, i.e. a purely physical phenomenon. Cenci et al. also confirmed that with application of ultrasound the reaction rates are over 200 times faster for the hydrolysis of benzoyl chloride in biphase CO₂/H₂O system (CO₂ 8 MPa, 30 °C).¹³⁴

Chatterjee et al.¹³⁵ studied cleavage of C–O bond of diphenyl ether by hydrolysis and hydrogenolysis in biphase CO₂/H₂O system in the presence of hydrogen (CO₂ 5–15 MPa, H₂ 0.5–4 MPa, 80 °C, 5 h). They used Rh@C catalyst and revealed combined synergetic effect of CO₂ and H₂O, which was evident from the product distribution (cyclohexane, cyclohexanol, dicyclohexyl ether) taking into account further hydrogenation (due to H₂ presence) of intermediates (benzene, phenol, cyclohexylphenyl ether).

3.10 Hydrolysis/condensation of organosilanes

We recently studied hydrolytic polycondensation of diethoxydimethylsilane in biphase CO₂/H₂O system.¹³ The reaction occurred at the temperatures in the range of 60–120 °C and the CO₂ saturating pressures of 43–77 MPa, t = 10–180 min. Complete conversion of the monomer was achieved even for as short exposure times as 10 min with the formation of cyclic and linear oligomers. It is interesting that the relative amount of cyclic products increased with the prolonging of the reaction duration but decreased with temperature rise. At the same time, the variation of the pressure did not affect apparently the composition of the polycondensation products. The possibility to obtain industrially important precursors of silicones using environmentally friendly chlorine-free synthetic method was thus demonstrated.

3.11 Oxidation reactions

Tsang et al.¹³⁶ studied aerobic oxidation of toluene and some other alkyl aromatic hydrocarbons in biphase CO₂/H₂O system. High degree of conversion of toluene to benzoic acid with high selectivity was observed using fluorous cobalt(II) species as a catalyst and NaBr as a promoter (15 MPa CO₂, 1 MPa O₂, 120 °C, 24 h). Partial oxidation of various alkyl aromatic hydrocarbons (p-xylene, ethylbenzene, 9,10-dihydroanthracene, 2-methyl anthracene) was also demonstrated in the same system.

He et al.¹³⁷ reported aerobic oxidation of secondary alcohols, benzhydrol, 1-phenylethanol and cyclohexanol, to corresponding ketones with Co(II)-based supported and unsupported catalysts and O₂ oxidant in the two types of biphase systems of CO₂/CO₂-expanded liquids: water and polyethylene glycol (PEG). The reaction occurred at the temperature of 70 °C, the partial O₂ pressure of 1.5 MPa and the partial CO₂ saturating pressure of 8.5 MPa, t = 9–24 h. Interesting, that very poor yield of product was obtained in water, whereas much higher yields were obtained in PEG, particularly for benzophenone from benzhydrol: 3% in water, 91% in PEG at the same conditions. Further examples on reactions in CO₂-expanded liquids may be found below in next sections.

Miao et al.¹³⁸ described oxidation of aliphatic, allylic, heterocyclic, and benzylic alcohols in biphase CO₂/H₂O system, see Fig. 36. Oxidation of both primary alcohols to aldehydes (partial oxidation) and secondary alcohols to ketones is an industrially important organic reaction. The authors explored 2,2,6,6-tetramethylpiperidine-1-oxyl (TEMPO) as a catalyst (metal-free catalysis). They emphasized that acidic condition (carbonic acid) is essential to the TEMPO-catalyzed oxidation of alcohols when simple, inexpensive, and biodegradable NaNO₂ is used as an activator of oxygen. An aerobic oxidation was performed, therefore oxygen was present in the reactor at the partial pressure of 1 MPa. The saturating CO₂ partial pressure was in range of 3–10 MPa, the temperature of 100 °C, t = 3–24 h. A variety of both primary and secondary alcohols were tested, with aryl and alkyl pendant groups. In many cases both complete conversions and quantitative yields were achieved, particularly at the saturating CO₂ pressure of 10 MPa. It is interesting that facile separation of the catalysts from the reaction mixture and its recycling was demonstrated. According to the results obtained all the tested primary benzylic alcohols, independently on the nature of substituents on the benzene ring, were converted into their corresponding aldehydes in high yields with excellent selectivity. Whereas the reactivity of the secondary benzylic alcohols was affected by the nature of the substituent groups probably due to the steric hindrance, thus, moderate yields were attained only at prolonged reaction times. Further, several aliphatic alcohols were also examined. The conversion degrees were mainly smaller for them being particularly poor for primary aliphatic alcohols with branched chains or secondary aliphatic alcohols. Formation of alkylcarbonic acids was detected from the alcohols in the presence of saturating CO₂.

Fig. 36 Oxidation of alcohols.¹³⁸

Liu et al.¹³⁹ performed selective oxybromination of electron-rich aromatics in biphase H₂O/CO₂ system in the presence of oxygen and bromine source (e.g., lithium bromide). High yields (up to 95%) with good regio-selectivity were achieved when cupric bromide was used as a catalyst at optimized reaction conditions (CO₂ 8 MPa, O₂ 1 MPa, 100 °C, 10 h). A variety of substrates were successfully converted, including aromatic ethers, sulfides and mesitylene.

4. Comparison with H₂O₂/CO₂ system, peroxycarbonic acid

Not only water may be saturated with CO₂ at high pressures in order to be converted into a new phase with unusual peculiar properties. Indeed, for example, if hydrogen peroxide is saturated with CO₂ in a high pressure vessel it is transformed into peroxycarbonic acid:

H₂O₂ + CO₂ → HOCOOOH

Peroxycarbonic acid is also unstable without saturating CO₂ and decomposes after decompression similarly to carbonic acid. Some generalizing reports, including accounts¹⁴⁰ and feature articles,¹⁴¹ were already published on application of peroxycarbonic acid as a reaction medium in organic synthesis. Therefore, we will only briefly mention most striking features of such reactions.

It is interesting to mention that H₂O₂ itself may be synthesized in carbon dioxide medium. Indeed, Hâncu et al.^142,143 produced hydrogen peroxide directly from O₂ and H₂ in CO₂ phase using CO₂-soluble Pd(II) or Pd(0) catalysts. Hydrogen peroxide is produced in quasi-organic CO₂ solvent and then it is stripped into water phase without any CO₂ “contamination” of the aqueous phase, which is typical for usual water/organic biphase systems. Hydrogen and oxygen dissolve well in CO₂ thus facilitating high rates of the reaction. Further, unlike usual organic solvents, CO₂ is fully oxidized and cannot be oxidized further by the H₂O₂. In spite of high interest,^144,145 this technology possibly requires further optimization from the standpoint of CO₂ negatively affecting the catalyst performance.^146,147

Nolen et al.¹⁴⁸ studied epoxidation of olefins in biphase CO₂/H₂O₂ system. Direct epoxidation of olefins without usage of any transition metals catalysts is an industrially promising approach of obtaining epoxides and their derivatives. Such intermediates are widely explored in industrial chemical syntheses of cosmetics, detergents, polymers, cross-linked materials of different nature, etc. The epoxidation of an alkene with a peroxyacid to give an oxirane is a typical example of Prilezhaev reaction. As the typical reaction to be studied, the authors selected epoxidation of cyclohexene. Here cyclohexene oxide and 1,2-cyclohexane diol are the two main products to be expected, but formation of some other byproducts is possible as well. The reaction was performed in aqueous H₂O₂ phase at the CO₂ saturating pressure of 12 MPa and the temperature of 40 °C. Previously, it was checked and proven that at such pressures cyclohexane is completely miscible with CO₂ phase. No epoxide was formed without CO₂, i.e. when nitrogen was used to generate high pressure instead in a control experiment. Yet, just in the presence of CO₂ the conversion achievable was rather poor. The authors revealed that addition of NaHCO₃ promotes further the epoxidation reaction (increased conversion was observed) apparently due to the increase of the effective concentration of peroxycarbonic acid species being formed. Indeed, much greater concentration of the conjugate base expectedly leaded to the improved conversion and yield as compared to week concentration of pure peroxycarbonic acid. With addition of NaHCO₃ the reaction occurred also without CO₂, i.e. when nitrogen was used to generate high pressure instead, but the epoxide selectivity is greatly improved in the presence of CO₂: the achievable selectivity was eight times higher. This trend in selectivity may be considered as an indicator of the existence of peroxycarbonic acid as far as in this case the reaction pathway apparently better follows the Prilezhaev reaction. For comparison the authors also successfully preformed epoxidation of water-soluble 3-cyclohexen-1-carboxylate sodium salt with very high epoxide yield of 89%. This observation proves that transfer of the hydrophobic alkene dissolved in the non-polar CO₂ phase to the catalytic aqueous phase with dissolved peroxycarbonic acid is a limiting stage of the reaction.

Rajagopalan et al.¹⁴⁹ demonstrated successful homogeneous oxidation (>85% epoxidation selectivity) of a variety of olefins (cyclohexene, styrene, 1-methylcyclohexene, 4-methylcyclohexene) in CO₂-expanded CH₃CN/H₂O₂/H₂O mixtures (4–8 MPa, 30–50 °C, 4–24 h). They found out that addition of pyridine dramatically promotes the epoxidation reaction. Beckman proposed direct synthesis of propylene oxide from propylene under influence of H₂O₂ generated in situ from H₂ and O₂ mixture in biphase CO₂/H₂O system.¹⁵⁰ Direct conversion of styrene into styrene carbonate catalyzed by sodium phosphotungstate/n-Bu₄NBr was studied in biphase CO₂/H₂O system in the presence of NaHCO₃, which markedly improved the conversion degree (CO₂ 2.4 MPa, 50 °C, 6 h).¹⁵¹ Zhao et al.¹⁵² reported successful epoxidation of styrene in the cetyltrimethylammonium bromide/H₂O/heptane/styrene/H₂O₂ emulsion system in the presence of compressed CO₂ (5.3 MPa, 40 °C, 7 h). Further, they¹⁵³ demonstrated direct preparation of propylene carbonate from propylene and CO₂ using quaternary ammonium heteropolyphosphatotungstate–quaternary ammonium halide catalytic system with anhydrous H₂O₂ as an oxidant through one-pot two-step process: (i) epoxidation (CO₂ 3 MPa (r.t.), 75 °C, 3.5 h) with (ii) subsequent cycloaddition (CO₂ 3 MPa (r.t.), 140 °C, 4 h).

Ganchegui & Leitner¹⁵⁴ studied oxybromination of phenol and aniline derivatives in biphase CO₂/H₂O₂ system, see Fig. 37. Brominated aromatic compounds are important as building blocks widely used in organic synthesis of fine chemicals. Firstly, the authors studied oxybromination of o-cresol, as a model substrate. The reaction was performed at the temperature of 40 °C in aqueous solution of H₂O₂ in the presence of NaBr, NaHCO₃ and at the saturating CO₂ pressure of 10–11 MPa, t = 4 h. In situ generation of peroxycarbonic acid from H₂O₂ and CO₂ was expected. Following ref. 148, the addition of NaHCO₃ facilitated formation of the sodium carboxylate of peroxycarbonic acid and as a result higher concentration of the overall peroxycarbonic species was also expected. Therefore, this approach also implies activation of hydrogen peroxide with bicarbonate. Finally, more than 90% conversion and 90% selectivity for the monobrominated product was achieved. For comparison, the conversion was only about 30% at the same conditions but without saturating CO₂.

Fig. 37 Oxybromination of o-cresol.¹⁵⁴

Secondly, the authors performed successful oxybromination of different phenol and aniline derivatives in the same reaction medium. In general, high conversions were achieved except for some strongly hydrophobic substrates, apparently, due to their poor solubility in water phase.

Karmee et al.¹⁵⁵ studied oxidation of sulfides to sulfones by in situ generated H₂O₂ in biphase CO₂/H₂O system. Conversion of thioanisole to methyl phenyl sulfone was used as a typical example. The one-pot process (H₂ 0.5–1.8 MPa, O₂ 0.5–0.6 MPa, H₂ + O₂ + CO₂ 13 MPa, 80 °C, 24 h) was catalyzed by two different nanoparticle catalysts: Pd@SiO₂ (formation of H₂O₂) and TiO₂ (oxidation of the sulfide to the sulfone).

Thus, comparing with previously considered biphase H₂O/CO₂ system (carbonic acid), one may notice that the focus of the reactions in the presence of peroxycarbonic acid is naturally shifted towards oxidation.

5. Comparison with biphase alcohol/CO₂ system, alkylcarbonic acid

Similarly to the situation with peroxycarbonic acid considered above, several reviews,^156,157 accounts^140,158 and feature articles¹⁴¹ were published on application of alkylcarbonic acids as a reaction media in organic synthesis. Therefore, here we will only briefly mention such reactions. Alkylcarbonic acids are formed when alcohols are saturated with CO₂ at high pressure:

ROH + CO₂ → ROCOOH

Alcohols saturated with CO₂ compose a subclass of broader class of gas-expanded solvents,^{156,158–163} which have attracted significant attention nowadays.^164–171 Indeed, mass transport is improved in such systems at sufficiently high gas content in the liquid phase. Their main properties, including dissolving power, may be tuned by pressure variation. Similarly to carbonic acid and peroxycarbonic acid, alkylcarbonic acids are unstable at normal pressure and therefore decompose after decompression and CO₂ release.

In order to prove directly formation of the acid, diazodiphenylmethane may serve as a reactive probe to trap the acid species for detection and comparison, see Fig. 38.¹⁷² Indeed, alcohols in the absence of an acid do not react significantly with this compound. Reaction rates of several alkylcarbonic acids with diazodiphenylmethane were studied and compared.¹⁷³ Thus, it was possible to estimate relative reaction rate based on acid strength. Different aliphatic and aromatic alcohols were selected as the source for alkylcarbonic acids. The authors detected the drop in polarity of the solvents with increasing pressure of saturating CO₂ thus hindering the proton transfer reaction, which becomes rate-determining step in corresponding reactions.

Fig. 38 Alkylcarbonic acid formation and reaction with diazodiphenylmethane to produce the ether and carbonate species.¹⁷²

Alkylcarbonic acids may be beneficial as a reaction media for certain reactions due to their lower polarity as compared to both carbonic and peroxycarbonic acids. Indeed, in many cases above we saw that the transfer of hydrophobic substrate to polar acid-containing aqueous phase limits the overall rate of the reaction. It may become particularly suppressed due to efficient extraction of the organic reactant from water phase by excessively dense pressurized CO₂. Alkylcarbonic acids are formed from alcohols, which initially already have significantly lower dielectric constants as compared to H₂O or H₂O₂. Carbon dioxide demonstrates much higher solubility in alcohols and even miscibility with some of them at sufficiently high pressures, i.e. it behaves significantly different as compared with its rather limited dissolution in polar aqueous media. Dissolution of large amount of non-polar molecules of carbon dioxide reduced significantly the dielectric constant of CO₂-saturated alcohols to be converted into alkylcarbonic acids. Therefore, such acids behave as low polar fluids with yet low pH values. Wen and Olesik determined that pH values of water/methanol mixtures saturated with CO₂ fell in the range of 4–5 and did not depend significantly on pressure variation in the range of 12–20 MPa.¹⁷⁴ Reduced polarity favors better dissolution of organic compounds. But, of course, reduced dielectric constant inevitably implies reduced ability of molecules to dissociate, therefore pH values cannot be smaller than the indicated boundary. Moreover, the reactions where proton transfer is essential should be also suppressed if the polarity of the medium is reduced so significantly. This possible obstacle should be taken into account.

Gohres et al.¹⁷⁵ investigated formation of alkylcarbonic acids and their dissociation constants in CO₂-saturated methanol, ethanol, and benzyl alcohol. They applied UV/vis spectroscopy with 2,6-dinitrophenol as an optical indicator. It was determined that methylcarbonic acid has a pK_a of 5.7. Ethylcarbonic acid is less acidic and benzyl carbonic acid is even lesser.

Liu et al.¹⁷⁶ repeated their previously described in ref. 111 experiments on partial reduction of nitroarenes to N-aryl hydroxylamines in the presence of ethanol. In particular, they focused on synthesis of N-(2-methylphenyl)hydroxylamine from o-nitrotoluene. The main features of the reaction occurrence were the same as described previously.¹¹¹ For example, when the CO₂ saturating pressure was fixed (P = 0.1 MPa), t = 1 h, and the temperature was varied, the dependence of both conversion and selectivity on temperature demonstrated non-monotonic dependence similarly to the one described before.¹¹¹ But the addition of ethanol increased the reaction selectivity towards N-(2-methylphenyl)hydroxylamine. They detected that at the increase of the ethanol portion, the selectivity towards this product was improved further, yet the reaction rate became low apparently due to decreased acidity of the reaction medium.

Bhanage et al.¹⁷⁷ performed synthesis of dimethyl carbonate and glycols from epoxides in CO₂-saturated methanol (8 MPa, 150 °C, 15 h + 4 h). Successful cycloaddition reaction of CO₂ with epoxides was also reported using relatively low pressure (1 MPa) and moderate temperature (85 °C) in green solvents (ethanol and propan-2-ol) forming alkylcarbonic acids.¹⁷⁸ Moderate to high yields and excellent selectivity of cyclic carbonates with vinyl or acrylate groups were achieved. Aoyagi et al.¹⁷⁹ performed similar reaction even at atmospheric pressure of CO₂ using phosphonium iodides as catalyst.

Xie et al.¹⁸⁰ studied formation of acetal from cyclohexanone in CO₂-saturated methanol and ethylene glycol. Performing the reaction in methanol during different times in the range of 2–10 h they determined pseudo-first-order rate constants of the reaction. The reaction rate constants increased when the temperature was increased in the range of 25–50 °C. But the dependence of the reaction rate constants on the pressure was non-monotonic: there was clear maximum localized in the range of 1–3 MPa. The exact localization of the maximum was a function of temperature: the higher the temperature, the higher the position of the maximum at the pressure scale. When the reaction was performed in CO₂-saturated ethylene glycol, already 2 mol% of dissolved CO₂ have a significant catalytic effect (solubility of CO₂ in ethylene glycol is worse as compared to its solubility in methanol). In both cases most probably alkylcarbonic acids formed in CO₂-saturated alcohols serve as catalysts.

Xie et al.¹⁸¹ further also analyzed hydrogenation of nitriles to amines in CO₂-saturated ethanol. Typically, nitrile hydrogenation is complicated by formation of secondary and tertiary amines. Therefore, certain protection is required to obtain primary amines. It is very interesting, that CO₂ may serve as such a protection agent for the primary amines. Indeed, when amines are saturated with CO₂ at high pressure, formation of carbamic acids occurs, similarly to the formation of carbonic acid, peroxycarbonic acid, alkylcarbonic acids, which were considered above. The carbamic acids may be further converted to carbamates. In any case, protection is achieved and selectivity towards the primary amines may be increased. The authors successfully performed hydrogenation of phenylacetonitrile and benzonitrile at the presence of nickel chloride and sodium borohydride in ethanol at the CO₂ saturating pressure of 3 MPa and at the temperature of 30 °C, t = 24 h. By means of the application of the protection with CO₂ indeed the yield and selectivity of the primary amines production was greatly increased.

Chamblee et al.¹⁸² studied hydrolysis of β-pinene in water–methanol mixtures saturated with CO₂. It was found out that most optimal was a 1 : 1 mixture (by volume) containing ∼3 mol% CO₂: from the viewpoint of yield and product distribution. The conversion of 71% was achieved for that system at the temperature of 75 °C for 24 h. The product distribution was: 14% hydrocarbons, 58% terpineol (mixture of isomers), 7% bicyclic alcohols, 19% methyl ether, which was rather favorable compared to typical literature examples. The most desirable product of this reaction is terpineol as far as is has a pleasant odor and is used in the flavor and fragrance industry. It is interesting to mention that without methanol—in pure water saturated with CO₂, i.e., carbonic acid—the conversion degree was 48% only, thus significantly smaller at otherwise the same conditions, but in the presence of methylcarbonic acid. This observation may be related to the fact that the substrate, β-pinene, is soluble in methanol, somewhat soluble in its mixtures with water, but insoluble in pure water.

Keating & Armstrong¹⁸³ realized Ugi reaction as one pot multi-component condensation of carbonyl compound, amine, and isocyanide with alkylcarbonic acid (t = 18–24 h). Several different alcohols (methanol, ethanol, allyl alcohol, 2-propanol, benzyl alcohol, 9-fluorenylmethanol) were saturated with CO₂ for formation of the corresponding acids. For heavy (solid) alcohols such as benzyl or 9-fluorenylmethyl, chloroform was additionally introduced as a non-reactive solvent. The product, N-(alkoxycarbonyl)aminoamide, was obtained with rather different yields (from very good to poor), but generally good for methanol, which was found to be the superior alcohol substrate. Similarly, Hulme et al.¹⁸⁴ also performed Ugi reaction as one pot multi-component condensation of aldehyde, amine, and isonitrile (isocyanide) with methylcarbonic acid (t = 18 h). The latter was generated in situ by saturation of methanol with CO₂. The reaction could proceed invariantly with a range of commercially available primary amines and isonitriles. The product, carbamate protected amino-amides, was obtained in good yields and could be further converted to cyclic ureas or hydantoins. These compounds are important for biomedical applications.¹⁸⁵

Weikel et al.¹⁸⁶ studied synthesis of methyl yellow and iodobenzene in CO₂-saturated methanol. The reaction was performed through formation of diazonium salts from aniline in the presence of NaNO₂, similarly to the synthetic route described in ref. 124, but with the difference that methylcarbonic acid was used as an activator instead of carbonic acid. The subsequent transformation of the diazonium salts to the corresponding products was achieved by their reaction with either N,N-dimethyl aniline or potassium iodide, respectively. For the synthesis of methyl yellow most optimal reaction temperature was 5 °C (t = 24 h). At these conditions yield of 72% methyl yellow and 97% conversion of aniline were achieved. By means of several control experiments (e.g., without CO₂, or with tetrahydrofuran instead of methanol, or without any solvent at all) it was demonstrated that methylcarbonic acid was essential for catalysis. For synthesis of iodobenzene the best yield of 72% was achieved at the highest temperature tested, 50 °C.

It is mentioned above, that Liu et al.¹³⁹ successfully performed selective oxybromination of aromatics in biphase H₂O/CO₂ system. They also studied copper(II)-catalyzed regioselective oxybromination of aromatics in CO₂/ethanol system.¹⁸⁷ Good to excellent yields as well as very good regioselectivity were achieved (CO₂ 1 MPa, O₂ 1 MPa, 100 °C, 10 h). The authors believed that formation of ethylcarbonic acid was crucial for the promotion of the reaction. Different aromatic substrates including differently substituted aromatic ethers and also sulfides were successfully converted. Similarly, they further demonstrated iron(III)-catalyzed aerobic oxidative iodination of electron-rich aromatics in CO₂/glycol system (CO₂ 1 MPa, O₂ 1 MPa, 100 °C, 10 h), where corresponding in situ formed alkylcarbonic acid played the same role.¹⁸⁸

Lin et al.¹⁸⁹ directly compared catalytic hydrogenolysis of benzylic compounds in biphase CO₂/H₂O system and methanol/CO₂ system in the presence of hydrogen (1 MPa). Using benzyl alcohol, 1-phenylethanol and 2-phenyl-2-propanol as the reagents, the corresponding conversion products: toluene, ethylbenzene and cumene, respectively, were obtained. The authors related the formation of the carbonic acid and methylcarbonic acid, respectively, (CO₂ saturation pressure 1 MPa, 50 °C) with the observed improved catalysis (Pd@C catalyst) of the hydrogenolysis. Therefore, both acids were comparatively favorable for the reaction.

Thus, comparing with previously considered biphase H₂O/CO₂ system (carbonic acid), one may notice that essentially the same reactions (in particular, cycloaddition, reduction, hydrogenation, hydrolysis, oxybromination) may be successfully realized in the presence of in situ formed alkylcarbonic acids.

Synthesis of polymers in alkylcarbonic acids was also shown to be promising.¹⁹⁰

6. Organic synthesis in biphase H₂O/CO₂ system at high temperatures

6.1 High-temperature water

Water itself is considered to be the most important resource for green chemistry reactions. At ambient temperatures it is a polar solvent suitable for dissolution of many inorganic compounds, it has neutral pH and low ion product. But physical properties of water change tremendously with increasing temperature above 150 °C. Percolated network of hydrogen bonds typical for liquid water at ambient conditions becomes disrupted at elevated temperatures. Although water molecules still form clusters, these formations remain local and non-percolated. Consequently, ion product is increased with corresponding decrease of dielectric constant. Water becomes less polar and suitable for dissolution of nonionic organic compounds. With ion product increased water not only can be a solvent for organic species, but can actively participate in both acid- and base-catalyzed reactions.

Chemical reactions in high-temperature water (HTW) have been studied extensively in the past few decades with several reviews on the matter published. First comprehensive review on HTW as a prospective medium for chemical reactions was given by Akiya and Savage:¹⁹¹ both physical properties and various types of reactions (ionic, polar, non-polar and radical) were thoroughly discussed. Brunner presented a two-part review on near critical and supercritical water^192,193 for hydrolysis and hydrothermal reactions as well as for supercritical water oxidation (SCWO) processes. More detailed analysis of HTW study was done by Brunner in ref. 194. Some other reviews on HTW should also be mentioned.^195,196 In the framework of the present review we should address the question: what new is introduced to the properties of HTW when it is additionally saturated with CO₂?

It is mentioned in ref. 192 that for acid-catalyzed hydrolysis reactions, addition of CO₂ to HTW may increase the selectivity and the yields of the products due to lower pH value of CO₂-enriched HTW. Several works on hydrolysis reactions in biomass processing were given as an example. We present in this part the sketch of CO₂–HTW system as applied for acid-catalyzed organic synthesis.

6.2 Properties of HTW saturated with CO₂

Thus, the HTW has a reduced dielectric constant, because at high temperatures hydrogen bonds between water molecules tend to break and consequently the medium becomes less polar. Therefore HTW shows an ability to dissolve some low-polarity species that are insoluble in usual water.¹⁹⁷

Hunter and Savage²⁰³ demonstrated that the dependence of the solubility constant K₀ (see eqn (6)) on the temperature exhibits a minimum at the temperature value around 150 °C, but the solubility of CO₂ in water further increases with temperature rise after this point. The solubility eventually achieves the larger values as compared to the room temperature (see Fig. 39). Nevertheless, the first apparent dissociation constant K^app₁ decreases dramatically with temperature (see Fig. 39) due to the decrease in dielectric constant of water.

Fig. 39 Solubility (K₀) and first apparent dissociation (K_a1) constants as a function of temperature. Reprinted from ref. 203, copyright© 2007, with permission from John Wiley and Sons.

Thus, the pH of the HTW saturated with CO₂ increases with temperature, see relevant discussion below. Indeed, as temperature increases, though the solubility of CO₂ in water increases (after 150 °C), but at the same time carbonic acid becomes less and less acidic. Nevertheless the pH of the HTW saturated with CO₂ is still lower as compared to the pH of the pure HTW without any saturating CO₂. The same is valid when water is transferred into sc state in the presence of CO₂.¹⁹⁸ Proton concentration naturally increases with the increase of the CO₂ partial pressure. At the same time it should be mentioned that the pH of pure HTW without any dissolved CO₂ also decreases with temperature from the neutral values at normal conditions to the value of 5.7 at 300 °C.¹⁹⁹

6.3 Acid-catalyzed reactions in HTW saturated with CO₂

To the best of our knowledge, first report on addition of CO₂ to HTW for catalyst enhancement was given by Alemán et al.²⁰⁰ The authors have shown that CO₂ promotes catalysis in methyl benzoate decarboxylation (Fig. 40) increasing conversion from 50 to 80%.

Fig. 40 Methyl benzoate decarboxylation reaction.²⁰⁰

First works specially focusing on the investigation of CO₂ influence on the reactions in HTW were reported by Hunter and Savage in ref. 201 and 202. The alkylation of p-cresol with tert-butyl alcohol to form 2-tert-butyl-4-methylphenol and the dehydration of cyclohexanol to form cyclohexene²⁰¹ as well as the inverse hydration of cyclohexene to form cyclohexanol²⁰² were carried out in HTW with and without CO₂, see Fig. 41.

Fig. 41 Upper raw: (left) cyclohexene molar yields in HTW and in CO₂-enriched HTW at T = 250 °C and 275 °C, (right) 2-tert-butyl-methylphenol molar yield in HTW and in CO₂-enriched HTW at T = 275 °C, reprinted from ref. 201, copyright© 2003, with permission from American Chemical Society. Lower raw: effect of CO₂ loading on cyclohexanol molar yield from cyclohexene hydration at T = 300 °C, reprinted from ref. 202, copyright© 2004, with permission from Elsevier.

In every case, yields of the products increased significantly with the addition of CO₂ as compared to pure HTW (Fig. 41). Authors attributed the observed enhancement of reaction rates to the pH decreasing with addition of CO₂. Calculations presented in ref. 201 and 203, allowed authors to plot the variation of pH with CO₂ partial pressure increase in the vessel containing biphase H₂O/CO₂ system at elevated temperatures. The thus obtained curves are presented in Fig. 42. It is clear from the curves that the best effect in terms of pH lowering by CO₂ addition can be achieved at moderately elevated temperatures of 150–200 °C. On the other hand, with the further temperature rise, although pH of CO₂-saturated HTW media tends towards neutral region, ionization constant of H₂O molecules is increasing with both polarity and viscosity of HTW decrease¹⁹¹ making water molecules themselves good hydrogen donors for acid-catalyzed organic reactions.

Fig. 42 Estimated pH of CO₂-enriched HTW as a function of CO₂ partial pressure, reprinted from ref. 203, copyright© 2007, with permission from John Wiley and Sons.

Miyazawa and Funazukuri exploited organic reaction rate acceleration with CO₂ addition into HTW to optimize polysaccharide hydrolysis reaction under hydrothermal conditions.²⁰⁴ As it can be seen from Fig. 43, monosaccharides yields increase significantly when CO₂-saturated HTW is used. Moreover, in the case of glucose yield from starch hydrolyses, this increase is almost proportional to the amount of CO₂ added to the vessel. Orozco et al.²⁰⁵ showed later that this batch process can be scaled successfully at least to a 250 ml vessel. If scaled further, the proposed polysaccharide hydrolysis route can be an environmentally friendly approach to biomass utilization process, such as xylol hydrolysis to furfural²⁰⁶ or inulin hydrolysis to 5-hydroxymethylfurfural.²⁰⁷ A full review on biomass utilization in the presence of CO₂, HTW–CO₂ system included, is given by Morais et al.²⁰⁸

Fig. 43 Yields of monosaccharides produced from polysaccharide hydrolyses with and without 0.2 g of carbon dioxide. Agar was degraded at 160 °C for 30 min. The other polysaccharides were degraded at 200 °C for 15 min. Reprinted from ref. 204, copyright© 2008, with permission from John Wiley and Sons.

Another important industrial application of CO₂/HTW system was proposed by Fan et al.²⁰⁹ Using model reaction of anthracene cracking under hydrothermal conditions they showed that HTW and CO₂ addition to the system not only can enhance hydrocracking process for producing liquid fuels from heavy oils, but also reduce the negative effect of nitrogen compounds in such process.²¹⁰

Concerning oxidation, Shang et al.²¹¹ described promoting effect of compressed CO₂ (1.3 MPa total pressure of O₂, N₂ and CO₂) on liquid phase oxidation of p-xylene in the presence of water and additional presence of acetic acid (186 °C). They used bromides of transition metals as catalysts. The authors ascribed promoting function of CO₂ to the formation of active peroxocarbonate from a synergistic interaction of CO₂ and O₂.

Investigating 1,4-butanediol dehydration to form tetrahydrofuran, Hunter et al. have shown that the reaction rate dependence on pH value is not a uniform one.²¹² While in acid or base pH regions the reaction rate constant grows with pH increase, it remains almost independent of pH in a pH region close to neutral values (see Fig. 44, left).

Fig. 44 Left: effect of pH and temperature on 1,4-butanediol dehydration in H₂O at elevated temperatures. Right, upper: acid-catalyzed mechanism for 1,4-butanediol dehydration with H⁺ serving as the catalyst, lower: acid-catalyzed mechanism for 1,4-butanediol dehydration with water serving as the catalyst. Reprinted from ref. 212, copyright© 2006, with permission from American Chemical Society.

To explain this non-uniform dependence, the authors proposed reaction mechanism that involves not only H⁺ acting as a catalyst, but also water molecules being a hydrogen donor (Fig. 44, right). Indeed, tetrahydrofuran formation from 1,4-butanediol is known to proceed in HTW without CO₂.²¹³ Further investigation of reaction rate enhancement with CO₂ addition to HTW done by Hunter and Savage²⁰³ confirms that the rate of the reactions that are of the first order in H⁺ increases significantly with CO₂ addition to HTW. For other reactions the expected increase is not so dramatic. Calculation were performed in ref. 203 to estimate the effect of CO₂ pressure in biphase H₂O/CO₂ system on the reaction rate of acid-catalyzed reactions that are of the first order in H⁺ and reactant, see Fig. 45.

Fig. 45 Effect of CO₂ addition, quantified in terms of the resulting increase in process operating pressure, on estimated rate increase factor for CO₂-enriched HTW, reprinted from ref. 203, copyright© 2007, with permission from John Wiley and Sons.

Results presented in Fig. 45 indicate that the most effective way to increase the reaction rate at moderate temperatures (150–200 °C) is to add CO₂ to HTW. On the other hand, with further temperature rise, the reaction mechanism with H₂O molecules serving as a proton donor becomes more favorable. Thus, both CO₂ addition and temperature rise can be considered as available means to promote reaction rates in HTW.

Proceeding with the research on the dehydration reactions of polyalcohols to form corresponding cyclic ethers, Yamaguchi et al. showed that the cyclizations of 2,5-hexanediol, 2,5-dimethyl-2,5-hexanediol, 1,4-butanediol, 1,4-pentanediol, 1,2,4-butanetriol and 1,2,5-pentanetriol are also promoted by addition of CO₂.^214–217 Acetol formation during 1,2,3 propanetriol dehydration in HTW is also enhanced with addition of CO₂.^214,218 Dehydration of D-sorbitol to form 1,4-anhydrosorbitol, 2,5-anhydrosorbitol and isosorbide was also promoted in CO₂–HTW media as opposed to pure HTW.¹⁰³

Sato et al. demonstrated for the first time, that synthesis in HTW with CO₂ addition can be performed not only in sealed reactor, but also in a continuous flow reactor system (Fig. 46).²¹⁹

Fig. 46 Schematic illustration of a flow reactor system for continuous dehydration of 1,4-butanediol in water with carbon dioxide used in ref. 219: (1) water, (2) carbon dioxide cylinder, (3) aqueous solution of 1,4-butandiol, (4–6) syringe pumps, (7) preheater, (8) oven, (9) tubular reactor, (10) condenser, (11) back pressure valve, and (12) gas–liquid separator, reprinted from ref. 219, copyright© 2015, with permission from Elsevier.

Proving that not only batch, but continuous synthesis can be performed is very important for acid-catalyzed reactions in CO₂–HTW system to be considered in the industry as a “green” alternative to conventional reaction routes using mineral acids as catalysts.

7. Conclusions

Possibility to generate self-neutralizing carbonic acid in biphase H₂O/CO₂ system opens new horizons for organic synthesis within the framework of green chemistry paradigm, particularly for acid-catalyzed reactions, while the typical problem of salt disposal is eliminated here completely.

Main physicochemical properties of carbonic acid (such as dielectric permittivity, density, pH) are to be tuned by the variations of external pressure or temperature. Carbonic acid is a not so week acid with the pK_a value of about 3.5, although only 0.2% of totally dissolved CO₂ molecules assemble into the H₂CO₃ molecules. The dependencies of solubility of water in CO₂ phase on pressure and CO₂ in the water phase on temperature demonstrate peculiar non-monotonic behavior. The molar fractions of CO₂ dissolved in water or water in CO₂ phase affect all the main properties of these media. Generally, the increase in pressure at constant temperature leads to the decrease of pH and dielectric constant of the water phase and to the increase of density, viscosity and dielectric constant of the CO₂ phase. The increase in temperature at constant pressure leads to the increase of dielectric constant of the CO₂ phase and to the decrease of density and dielectric constant of the water phase. In particular, the decrease in dielectric constant of water saturated with CO₂ in comparison with pure water can improve the solubility of some organic species in it. Similarly, the increase in dielectric constant of CO₂ in a contact with water under pressure in comparison with pure CO₂ can improve the solubility of some polar species in it.

Formation of other metastable self-neutralizing acids, such as peroxycarbonic acid, alkylcarbonic acids, carbamic acids, may also be beneficial for some reaction routes in corresponding biphase systems. The composition, acidity, density, viscosity, dielectric constant of one or both phases in such biphase systems may also be conveniently adjusted by means of temperature and pressure variation.

Hydrolysis and oxidation reactions seem to be particularly interesting to be performed in the presence of peroxycarbonic acid – but, however, just in the presence of “simple” carbonic acid also – as far as the non-polar solvent of the biphase system, CO₂, is already in a completely oxidized stable state in spite of its non-polar nature, which is a rather atypical situation for biphase systems.

Yet, the question about the most optimal pressure and temperature regions for the organic synthesis does not still have direct and straightforward answer. Only certain tendencies may be outlined. Firstly, when pressure is fixed, and temperature is varied, there is apparently a “dead valley” region located around 60 °C (the exact location may be a function of the selected pressure), where acid-catalyzed reactions in biphase H₂O/CO₂ system in general proceed particularly ineffectively. In such cases some shift of the temperature towards either higher or lower values may be recommended. The existence of this “dead valley” region is apparently due to competition between thermal activation of the reaction and reduced acidity of carbonic acid solution in the biphase H₂O/CO₂ system at temperature rise. Secondly, when temperature is fixed, and pressure is varied, for some reactions there is a maximum of reaction efficiency located at the value of several MPa (the exact location is a function of the selected temperature). At further pressure increase the pressurized dense CO₂ efficiently extracts the organic reagent from water phase, therefore preventing its contact with a catalyst and the reaction. As a result, rather often the reaction rate is reduced at excessive CO₂ pressures.

This ability of dense CO₂ to affect significantly the miscibility of water and organics dissolved in it as well as to induce their phase separation²²⁰ may be not only detrimental but beneficial either for certain reaction systems. Indeed, the selective and pressure-tunable solubility of products, reagents, intermediates or/and catalysts makes the biphase system to be a convenient polar/nonpolar system without any cross-contamination problem, which may ensure simplified product removal, catalyst recycling or product replenishment as well as unique selectivity achievable. This convenient polar/nonpolar phase may be explored in any appropriate chemical reactions requiring usage of biphase systems in general, i.e., not only acid-catalyzed ones.

Another question to be addressed is whether high pressure of CO₂ in H₂O/CO₂ or similar biphase systems is really required at all? Indeed, we saw that in some cases just simple bubbling of water with CO₂ at atmospheric pressure already provided the pH values required for noticeable reaction rates. Nevertheless, the general experimental tendency is that the pressure increase above atmospheric one always significantly facilitates the course of all the reactions studied so far. Therefore, taking into account the substantial recent progress with high pressure equipment available nowadays (easy-to-use pressure generators, sophisticated autoclaves), the shift towards pressures higher than atmospheric one, at least, seems to be strongly recommended. Such pressures applied may ensure rates, yields and selectivity achievable being sufficiently high as to make the reactions acceptable for industrial implementation.

Among other options concerning biphase H₂O/CO₂ system one can mention possibility to form emulsions or to apply ultrasound agitation in order to speed up the reaction rate. The emulsions are generally broken at decompression, which facilitate product/catalysts separation and recovery. Besides, CO₂ may introduce chemical protection to certain groups, which bring on additional advantageous selectivity. One can also mention that usual liquid acid may be added to the system in order to catalytically promote the reaction, if necessary.²²¹ But there is a very interesting example of usage of solid superacidic hydrated Nafion film both as a catalyst and a source of water for hydration reactions,²²² similar to the ones typically performed in biphase H₂O/CO₂ system.²²³ Important benefit of the usage of a solid Nafion film as a catalytic promoter is that it is extremely easy in handling and recovery.

High temperature water is also a promising media with unusual properties for “green” acid- and base-catalyzed reactions. The saturation of HTW with CO₂ under pressure additionally lowers its pH as well as the dielectric constant as compared to pure HTW. Thus, the CO₂ addition to HTW promotes acid-catalyzed reactions that are of the first order in H⁺. Theoretically, for such reactions the reaction rate can be increased by two orders of magnitude with CO₂ pressure penalty of only 0.2–1 MPa. However, with further increase of temperature above 200 °C the effectiveness of CO₂ addition to promote the reaction lessens significantly. Apparently, at temperatures of 250–300 °C H₂O molecules themselves play a key role in catalysis being a direct hydrogen donor.

Acknowledgements

This work was financially supported by Russian Science Foundation (project no. 14-23-00231).

References

1. E. J. Beckman, J. Supercrit. Fluids, 2004, 28, 121–191.
2. P. G. Jessop, T. Ikariya and R. Noyori, Chem. Rev., 1999, 99, 475–493.
3. J. A. Darr and M. Poliakoff, Chem. Rev., 1999, 99, 495–541.
4. A. Baiker, Chem. Rev., 1999, 99, 453–473.
5. J. L. Kendall, D. A. Canelas, J. L. Young and J. M. DeSimone, Chem. Rev., 1999, 99, 543–563.
6. P. E. Savage, Chem. Rev., 1999, 99, 603–621.
7. W. Leitner, Acc. Chem. Res., 2002, 35, 746–756.
8. P. G. Jessop, J. Supercrit. Fluids, 2006, 38, 211–231.
9. E. Ramsey, Q. Sun, Z. Zhang, C. Zhang and W. Gou, J. Environ. Sci., 2009, 21, 720–726.
10. I. Omae, Coord. Chem. Rev., 2012, 256, 1384–1405.
11. D. Yu, S. P. Teong and Y. Zhang, Coord. Chem. Rev., 2015, 293–294, 279–291.
12. B. Walsh, J. R. Hyde, P. Licence and M. Poliakoff, Green Chem., 2005, 7, 456–463.
13. A. A. Kalinina, I. V. Elmanovich, M. N. Temnikov, M. A. Pigaleva, A. S. Zhiltsov, M. O. Gallyamov and A. M. Muzafarov, RSC Adv., 2015, 5, 5664–5666.
14. S. Wu, H. Fan, Y. Cheng, Q. Wang and B. Han, Progr. Chem., 2010, 22(7), 1286–1294.
15. M. Arai, S.-I. Fujita and M. Shirai, J. Supercrit. Fluids, 2009, 47, 351–356.
16. G. R. Akien and M. Poliakoff, Green Chem., 2009, 11, 1083–1100.
17. A. Z. Fadhel, P. Pollet, C. L. Liotta and C. A. Eckert, Molecules, 2010, 15, 8400–8424.
18. Y. Medina-Gonzalez, S. Camy and J.-S. Condoret, ACS Sustainable Chem. Eng., 2014, 2, 2623–2636.
19. J. J. de Vries, in Environmental Isotopes in the Hydrological Cycle: Principles and Applications, ed. W. G. Mook, INEA/UNESCO, Paris, 2000, vol. 9, pp. 143–165.
20. R. K. Lam, A. H. England, J. W. Smith, A. M. Rizzuto, O. Shih, D. Prendergast and R. J. Saykally, Chem. Phys. Lett., 2015, 633, 214–217.
21. R. K. Lam, A. H. England, A. T. Sheardy, O. Shih, J. W. Smith, A. M. Rizzuto, D. Prendergast and R. J. Saykally, Chem. Phys. Lett., 2014, 614, 282–286.
22. W. Stumm and J. J. Morgan, Aquatic Chemistry: Chemical Equilibria And Rates In Natural Waters, John Wiley & Sons, Inc., New York, 1996.
23. W. P. Jencks and R. A. Altura, J. Chem. Educ., 1988, 65, 770–771.
24. S. K. Reddy and S. Balasubramanian, Chem. Commun., 2014, 50, 503–514.
25. D. M. Kern, J. Chem. Educ., 1960, 37, 14–23.
26. A. J. Read, J. Solution Chem., 1975, 4, 53–70.
27. K. Adamczyk, M. Prémont-Schwarz, D. Pines, E. Pines and E. T. J. Nibbering, Science, 2009, 326, 1690–1694.
28. W. Henry, Philos. Trans. R. Soc. London, 1803, 93, 29–42 , 274–276.
29. N. Spycher, K. Pruess and J. Ennis-King, Geochim. Cosmochim. Acta, 2003, 67, 3015–3031.
30. R. Wiebe, Chem. Rev., 1941, 29, 475–481.
31. E. P. Bartlett, J. Am. Chem. Soc., 1927, 49(1), 65–78.
32. H. Teng, A. Yamasaki, M.-K. Chun and H. Lee, J. Chem. Thermodyn., 1997, 29, 1301–1310.
33. C. Peng, J. P. Crawshaw, G. C. Maitland, J. P. M. Trusler and D. Vega-Maza, J. Supercrit. Fluids, 2013, 82, 129–137.
34. K. L. Toews, R. M. Shroll, C. M. Wai and N. G. Smart, Anal. Chem., 1995, 67, 4040–4043.
35. C. Roosen, M. Ansorge-Schumacher, T. Mang, W. Leitner and L. Greiner, Green Chem., 2007, 9, 455–458.
36. M. A. Khokhlova, M. O. Gallyamov and A. R. Khokhlov, Colloid Polym. Sci., 2012, 290(15), 1471–1480.
37. M. O. Gallyamov, I. S. Chaschin, M. A. Khokhlova, T. E. Grigorev, N. P. Bakuleva, I. G. Lyutova, J. E. Kondratenko, G. A. Badun, M. G. Chernysheva and A. R. Khokhlov, Mater. Sci. Eng., C, 2014, 37, 127–140.
38. M. A. Pigaleva, I. V. Portnov, A. A. Rudov, I. V. Blagodatskikh, T. E. Grigoriev, M. O. Gallyamov and I. I. Potemkin, Macromolecules, 2014, 47, 5749–5758.
39. M. A. Pigaleva, M. V. Bulat, G. N. Bondarenko, S. S. Abramchuk, T. V. Laptinskaya, M. O. Gallyamov, I. P. Beletskaya and M. Möller, ACS Macro Lett., 2015, 4, 661–664.
40. M. B. King, A. Mubarak, J. D. Kim and T. R. Bott, J. Supercrit. Fluids, 1992, 5, 296–302.
41. A. Hebach, A. Oberhof and N. Dahmen, J. Chem. Eng. Data, 2004, 49, 950–953.
42. M. McBride-Wright, G. C. Maitland and J. P. M. Trusler, J. Chem. Eng. Data, 2015, 60, 171–180.
43. R. Yaginuma, Y. Sato, D. Kodama, H. Tanaka and M. Kato, J. Jpn. Inst. Energy, 2000, 79, 144–146.
44. K. M. Sabil, G.-J. Witkamp and C. J. Peters, J. Chem. Thermodyn., 2010, 42, 8–16.
45. K. M. Sabil, G.-J. Witkamp and C. J. Peters, J. Chem. Eng. Data, 2010, 55(2), 813–818.
46. J. A. Hyatt, J. Org. Chem., 1984, 49(26), 5097–5101.
47. A. Michels and L. Kleerekoper, Physica, 1939, 6, 586–590.
48. A. Michels, J. de Boer and A. Bijl, Physica, 1937, 4, 981–994.
49. A. Michels and C. Michels, Philos. Trans. R. Soc., A, 1933, 231, 409–434.
50. A. Catenaccio, Y. Daruich and C. Magallanes, Chem. Phys. Lett., 2003, 367, 669–671.
51. W. B. Floriano and M. A. C. Nascimento, Braz. J. Phys., 2004, 34, 38–41.
52. U. Kaatze, J. Solution Chem., 1997, 26, 1049–1112.
53. P. Wang and A. Anderko, Fluid Phase Equilib., 2001, 186, 103–122.
54. J. G. Kirkwood, J. Chem. Phys., 1939, 7, 911–919.
55. R. M. Fuoss and J. G. Kirkwood, J. Am. Chem. Soc., 1941, 63, 385–394.
56. A. I. Cooper, J. D. Londono, G. Wignall, J. B. McClain, E. T. Samulski, J. S. Lin, A. Dobrynin, M. Rubinstein, A. L. C. Burke, J. M. J. Fréchet and J. M. DeSimone, Nature, 1997, 389, 368–371.
57. L. Yang, Q. Pan and G. L. Rempel, Catal. Today, 2013, 207, 153–161.
58. Y. Himeda, N. Onozawa-Komatsuzaki, H. Sugihara and K. Kasuga, Organometallics, 2007, 26(3), 702–712.
59. H. Cheng, X. Meng, R. Liu, Y. Hao, Y. Yu, S. Cai and F. Zhao, Green Chem., 2009, 11, 1227–1231.
60. J. Sun, J. Ren, S. Zhang and W. Cheng, Tetrahedron Lett., 2009, 50, 423–426.
61. A. Dibenedetto and A. Angelini, in Advances in Inorganic Chemistry, ed. M. Aresta and R. van Eldik, Elsevier Inc., Amsterdam, 2014, vol. 66, suppl. 2, pp. 25–81.
62. D.-H. Lan, F.-M. Yang, S.-L. Luo, C.-T. Au and S.-F. Yin, Carbon, 2014, 73, 351–360.
63. R.-J. Wei, X.-H. Zhang, B.-Y. Du, Z.-Q. Fan and G.-R. Qi, J. Mol. Catal. A: Chem., 2013, 379, 38–45.
64. J. Ma, J. Song, H. Liu, J. Liu, Z. Zhang, T. Jiang, H. Fan and B. Han, Green Chem., 2012, 14, 1743–1748.
65. J. Ma, J. Liu, Z. Zhang and B. Han, Green Chem., 2012, 14, 2410–2420.
66. P. Mazo and L. Rios, J. Am. Oil Chem. Soc., 2013, 90, 725–730.
67. N. Eghbali and C.-J. Li, Green Chem., 2007, 9, 213–215.
68. B. Gabriele, G. Salerno, M. Costa and G. P. Chiusoli, Chem. Commun., 1999, 1381–1382.
69. G. P. Chiusoli, M. Costa, L. Cucchia, B. Gabriele, G. Salerno and L. Veltri, J. Mol. Catal. A: Chem., 2003, 204–205, 133–142.
70. B. Gabriele, L. Veltri, G. Salerno, M. Costa and G. P. Chiusoli, Eur. J. Org. Chem., 2003, 1722–1728.
71. J.-H. Li and Y.-X. Xie, Synth. Commun., 2004, 34(9), 1737–1743.
72. J.-H. Li and Y.-X. Xie, Chin. J. Chem., 2004, 22(12), 1421–1424.
73. J. P. Hanrahan, K. J. Ziegler, J. P. Galvin and J. D. Holmes, Langmuir, 2004, 20(11), 4386–4390.
74. M. T. Timko, A. J. Allen, R. L. Danheiser, J. I. Steinfeld, K. A. Smith and J. W. Tester, Ind. Eng. Chem. Res., 2006, 45(5), 1594–1603.
75. S. Minakata, I. Sasaki and T. Ide, Angew. Chem., Int. Ed., 2010, 49, 1309–1311.
76. J. Ma, B. Han, J. Song, J. Hu, W. Lu, D. Yang, Z. Zhang, T. Jiang and M. Hou, Green Chem., 2013, 15, 1485–1489.
77. W. Lu, J. Ma, J. Hu, Z. Zhang, C. Wu and B. Han, RSC Adv., 2014, 4, 50993–50997.
78. M. McCarthy, H. Stemmer and W. Leitner, Green Chem., 2002, 4, 501–504.
79. L. Schmid, M. S. Schneider, D. Engel and A. Baiker, Catal. Lett., 2003, 88(3–4), 105–113.
80. W. Cao, H. Wang, X. Gao, J. Zhang and J. Zhang, Chin. J. Catal., 2004, 25(7), 556–560.
81. J. P. Hallett, J. W. Ford, R. S. Jones, P. Pollet, C. A. Thomas, C. L. Liotta and C. A. Eckert, Ind. Eng. Chem. Res., 2008, 47, 2585–2589.
82. X. Xie, J. S. Brown, P. J. Joseph, C. L. Liotta and C. A. Eckert, Chem. Commun., 2002, 1156–1157.
83. C. D. Ablan, J. P. Hallett, K. N. West, R. S. Jones, C. A. Eckert, C. L. Liotta and P. G. Jessop, Chem. Commun., 2003, 2972–2973.
84. S. M. Mercer, T. Robert, D. V. Dixon and P. G. Jessop, Catal. Sci. Technol., 2012, 2, 1315–1318.
85. M. J. Lazzaroni, D. Bush, R. Jones, J. P. Hallett, C. L. Liotta and C. A. Eckert, Fluid Phase Equilib., 2004, 224, 143–154.
86. K. Burgemeister, G. Francio, H. Hugl and W. Leitner, Chem. Commun., 2005, 6026–6028.
87. K. Burgemeister, G. Franciò, V. H. Gego, L. Greiner, H. Hugl and W. Leitner, Chem.–Eur. J., 2007, 13, 2798–2804.
88. G. B. Jacobson, C. T. Lee, K. P. Johnston and W. Tumas, J. Am. Chem. Soc., 1999, 121, 11902–11903.
89. M. Ohde, H. Ohde and C. M. Wai, Chem. Commun., 2002, 2388–2389.
90. M. Ohde, H. Ohde and C. M. Wai, in Supercritical Carbon Dioxide, ed. A. Gopalan, et al., ACS Symposium Series, American Chemical Society, Washington, DC, 2003, vol. 27, pp. 419–429.
91. M. Ohde, H. Ohde and C. M. Wai, Langmuir, 2005, 21, 1738–1744.
92. R. Liu, C. Wu, Q. Wang, J. Ming, Y. Hao, Y. Yu and F. Zhao, Green Chem., 2009, 11, 979–985.
93. R. Liu, F. Zhao, S.-I. Fujita and M. Arai, Appl. Catal., A, 2007, 316, 127–133.
94. R. Ma, L.-N. He, Q.-W. Song, Y.-B. Zhou and K.-X. Liu, RSC Adv., 2014, 4, 11867–11871.
95. H.-F. Jiang and X.-Z. Huang, J. Supercrit. Fluids, 2007, 43, 291–294.
96. B. M. Bhanage, Y. Ikushima, M. Shirai and M. Arai, Chem. Commun., 1999, 1277–1278.
97. S.-I. Fujita, S. Akihara and M. Arai, J. Mol. Catal. A: Chem., 2006, 249, 223–229.
98. X.-Z. Huang and H.-F. Jiang, Chem. Res. Chin. Univ., 2008, 24(5), 658–660.
99. N. Hiyoshi, O. Sato, A. Yamaguchi and M. Shirai, Chem. Commun., 2011, 47, 11546–11548.
100. M. Chatterjee, T. Ishizakaa and H. Kawanami, Green Chem., 2014, 16, 1543–1551.
101. F. Liu, M. Audemar, K. de Oliveira Vigier, J.-M. Clacens, F. de Campo and F. Jérôme, ChemSusChem, 2014, 7, 2089–2093.
102. G. Li, H. Jiang and J. Li, Green Chem., 2001, 3, 250–251.
103. M. Shirai, O. Sato, N. Hiyoshi and A. Yamaguchi, J. Chem. Sci., 2014, 126, 395–401.
104. R. Ma, A.-H. Liu, C.-B. Huang, X.-D. Li and L.-N. He, Green Chem., 2013, 15, 1274–1279.
105. H. Yoshida, Y. Wang, S. Narisawa, S.-I. Fujita, R. Liu and M. Arai, Appl. Catal., A, 2013, 456, 215–222.
106. H. Cheng, X. Meng, C. Wu, X. Shan, Y. Yu and F. Zhao, J. Mol. Catal. A: Chem., 2013, 379, 72–79.
107. H. Yoshida, S. Narisawa, S.-I. Fujita and M. Arai, J. Mol. Catal. A: Chem., 2013, 379, 80–85.
108. A. Bhosale, H. Yoshida, S.-I. Fujita and M. Arai, Green Chem., 2015, 17, 1299–1307.
109. H. Yoshida, S.-I. Fujita and M. Arai, Catal. Today, 2015, 247, 170–176.
110. H.-F. Jiang and Y.-S. Dong, Chin. J. Chem., 2008, 26, 1407–1410.
111. S. Liu, Y. Wang, J. Jiang and Z. Jin, Green Chem., 2009, 11, 1397–1400.
112. S. Liu, Y. Wang, X. Yang and J. Jiang, Res. Chem. Intermed., 2012, 38, 2471–2478.
113. G. Gao, Y. Tao and J. Jiang, Green Chem., 2008, 10, 439–441.
114. H. Cheng, X. Meng, Y. Yu and F. Zhao, Appl. Catal., A, 2013, 455, 8–15.
115. J. Li, Y. Xie, H. Jiang and M. Chen, Green Chem., 2002, 4, 424–425.
116. J.-H. Li, Y.-X. Xie and D.-L. Yin, J. Org. Chem., 2003, 68, 9867–9869.
117. J.-H. Li and Y.-X. Xie, Chin. J. Chem., 2004, 22(9), 966–970.
118. F. Zhang, W. Chai and H.-X. Le, Chin. J. Chem., 2008, 26(1), 25–29.
119. B. M. Bhanage, Y. Ikushima, M. Shirai and M. Arai, Tetrahedron Lett., 1999, 40, 6427–6430.
120. B. M. Bhanage, S.-I. Fujita and M. Arai, J. Organomet. Chem., 2003, 687, 211–218.
121. N. Ye, X. Zhang and W. Shang, Ind. Eng. Chem. Res., 2014, 53, 14158–14165.
122. S. Liu, Y. Hao and J. Jiang, Ind. Eng. Chem. Res., 2014, 53, 8372–8375.
123. B. B. Zhang, J. L. Song, J. Ma, W. T. Wang, P. Zhang, T. Jiang and B.-X. Han, Sci. China: Chem., 2013, 56(10), 1436–1439.
124. P. Tundo, A. Loris and M. Selva, Green Chem., 2007, 9, 777–779.
125. G. B. Jacobson, C. T. Lee and K. P. Johnston, J. Org. Chem., 1999, 64, 1201–1206.
126. G. B. Jacobson, C. T. Lee, S. R. P. daRocha and K. P. Johnston, J. Org. Chem., 1999, 64, 1207–1210.
127. L. A. Blanchard and J. F. Brennecke, Green Chem., 2001, 3, 17–19.
128. R. Zhang, L. Guo, J. Chen, H. Gan, B. Song, W. Zhu, L. Hua and Z. Hou, ACS Sustainable Chem. Eng., 2014, 2, 1147–1154.
129. S. M. Cenci, L. R. Cox and G. A. Leeke, ACS Sustainable Chem. Eng., 2014, 2, 1280–1288.
130. J.-H. Li, Y.-X. Xie and D.-L. Yin, Green Chem., 2002, 4, 505–506.
131. C. M. Rayner, Org. Process Res. Dev., 2007, 11, 121–132.
132. M. T. Timko, J. M. Diffendal, J. W. Tester, K. A. Smith, W. A. Peters, R. L. Danheiser and J. I. Steinfeld, J. Phys. Chem. A, 2003, 107(29), 5503–5507.
133. M. T. Timko, K. A. Smith, R. L. Danheiser, J. I. Steinfeld and J. W. Tester, AIChE J., 2006, 52, 1127–1141.
134. S. M. Cenci, L. R. Cox and G. A. Leeke, Ultrason. Sonochem., 2014, 21, 401–408.
135. M. Chatterjee, A. Chatterjee, T. Ishizaka and H. Kawanami, Catal. Sci. Technol., 2015, 5, 1532–1539.
136. S. C. Tsang, J. Zhu and K. M. K. Yu, Catal. Lett., 2006, 106(3–4), 123–126.
137. J. He, T. Wu, T. Jiang, X. Zhou, B. Hu and B. Han, Catal. Commun., 2008, 9, 2239–2243.
138. C.-X. Miao, L.-N. He, J.-L. Wang and F. Wu, J. Org. Chem., 2010, 75, 257–260.
139. A.-H. Liu, R. Ma, M. Zhang and L.-N. He, Catal. Today, 2012, 194, 38–43.
140. J. P. Hallett, P. Pollet, C. L. Liotta and C. A. Eckert, Acc. Chem. Res., 2008, 41(3), 458–467.
141. C. A. Eckert, C. L. Liotta, D. Bush, J. S. Brown and J. P. Hallett, J. Phys. Chem. B, 2004, 108, 18108–18118.
142. D. Hâncu, J. Green and E. J. Beckman, Acc. Chem. Res., 2002, 35, 757–764.
143. D. Hâncu, J. Green and E. J. Beckman, Ind. Eng. Chem. Res., 2002, 41, 4466–4474.
144. S. Abate, S. Perathoner and G. Centi, Top. Catal., 2011, 54, 718–728.
145. T. M. Rueda, J. G. Serna and M. J. C. Alonso, J. Supercrit. Fluids, 2012, 61, 119–125.
146. A. Pashkova, L. Greiner, U. Krtschil, C. Hofmann and R. Zapf, Appl. Catal., A, 2013, 464–465, 281–287.
147. A. Pashkova and R. Dittmeyer, Catal. Today, 2015, 248, 128–137.
148. H. A. Nolen, J. Lu, J. S. Brown, P. Pollet, B. C. Eason, K. N. Griffith, R. Gläser, D. Bush, D. R. Lamb, C. L. Liotta, C. A. Eckert, G. F. Thiele and K. A. Bartels, Ind. Eng. Chem. Res., 2002, 41, 316–323.
149. B. Rajagopalan, M. Wei, G. T. Musie, B. Subramaniam and D. H. Busch, Ind. Eng. Chem. Res., 2003, 42(25), 6505–6510.
150. E. J. Beckman, Green Chem., 2003, 5, 332–336.
151. J.-L. Wang, J.-Q. Wang, L.-N. He, X.-Y. Dou and F. Wu, Green Chem., 2008, 10, 1218–1223.
152. Y. Zhao, J. Zhang, B. Han, S. Hu and W. Li, Green Chem., 2010, 12, 452–457.
153. G. Zhao, Y. Zhang, H. Zhang, J. Li and S. Gao, J. Energy Chem., 2015, 24, 353–358.
154. B. Ganchegui and W. Leitner, Green Chem., 2007, 9, 26–29.
155. S. K. Karmee, L. Greiner, A. Kraynov, T. E. Müller, B. Niemeijer and W. Leitner, Chem. Commun., 2010, 46, 6705–6707.
156. G. Musie, M. Wei, B. Subramaniam and D. H. Busch, Coord. Chem. Rev., 2001, 219–221, 789–820.
157. E. A. Karakhanov and A. L. Maksimov, Russ. J. Gen. Chem., 2009, 79(6), 1370–1383.
158. J. P. Hallett, C. L. Kitchens, R. Hernandez, C. L. Liotta and C. A. Eckert, Acc. Chem. Res., 2006, 39, 531–538.
159. B. Subramaniam, C. J. Lyon and V. Arunajatesan, Appl. Catal., B, 2002, 37, 279–292.
160. M. Wei, G. T. Musie, D. H. Busch and B. Subramaniam, J. Am. Chem. Soc., 2002, 124, 2513–2517.
161. B. Kerler, R. E. Robinson, A. S. Borovik and B. Subramaniam, Appl. Catal., B, 2004, 49, 91–98.
162. P. G. Jessop and B. Subramaniam, Chem. Rev., 2007, 107, 2666–2694.
163. B. Subramaniam, R. V. Chaudhari, A. S. Chaudhari, G. R. Akien and Z. Xie, Chem. Eng. Sci., 2014, 115, 3–18.
164. D. J. Heldebrant and P. G. Jessop, J. Am. Chem. Soc., 2003, 125(19), 5600–5601.
165. R. A. Sheldon, Green Chem., 2005, 7, 267–278.
166. P. G. Jessop, D. J. Heldebrant, X. Li, C. A. Eckert and C. L. Liotta, Nature, 2005, 436, 1102.
167. D. J. Heldebrant, P. K. Koech, M. T. C. Ang, C. Liang, J. E. Rainbolt, C. R. Yonkera and P. G. Jessop, Green Chem., 2010, 12, 713–721.
168. P. Pollet, R. J. Hart, C. A. Eckert and C. L. Liotta, Acc. Chem. Res., 2010, 43(9), 1237–1245.
169. A. Z. Fadhel, P. Pollet, C. L. Liotta and C. A. Eckert, Annu. Rev. Chem. Biomol. Eng., 2011, 2, 189–210.
170. X. Han and M. Poliakoff, Chem. Soc. Rev., 2012, 41, 1428–1436.
171. P. Pollet, E. A. Davey, E. E. Ureña-Benavides, C. A. Eckert and C. L. Liotta, Green Chem., 2014, 16, 1034–1055.
172. K. N. West, C. Wheeler, J. P. McCarney, K. N. Griffith, D. Bush, C. L. Liotta and C. A. Eckert, J. Phys. Chem. A, 2001, 105, 3947–3948.
173. R. R. Weikel, J. P. Hallett, C. L. Liotta and C. A. Eckert, Top. Catal., 2006, 37(2–4), 75–80.
174. D. Wen and S. V. Olesik, Anal. Chem., 2000, 72, 475–480.
175. J. L. Gohres, A. T. Marin, J. Lu, C. L. Liotta and C. A. Eckert, Ind. Eng. Chem. Res., 2009, 48(3), 1302–1306.
176. S. J. Liu, Y. H. Wang, Y. P. Hao and J. Y. Jiang, Chin. Chem. Lett., 2011, 22, 221–224.
177. B. M. Bhanage, S.-I. Fujita, Y. Ikushima and M. Arai, Appl. Catal., A, 2001, 219, 259–266.
178. A. J. R. Amaral, J. F. J. Coelho and A. C. Serra, Tetrahedron Lett., 2013, 54, 5518–5522.
179. N. Aoyagi, Y. Furusho and T. Endo, Tetrahedron Lett., 2013, 54, 7031–7034.
180. X. Xie, C. L. Liotta and C. A. Eckert, Ind. Eng. Chem. Res., 2004, 43, 2605–2609.
181. X. Xie, C. L. Liotta and C. A. Eckert, Ind. Eng. Chem. Res., 2004, 43, 7907–7911.
182. T. S. Chamblee, R. R. Weikel, S. A. Nolen, C. L. Liotta and C. A. Eckert, Green Chem., 2004, 6, 382–386.
183. T. A. Keating and R. W. Armstrong, J. Org. Chem., 1998, 63, 867–871.
184. C. Hulme, L. Ma, J. J. Romano, G. Morton, S.-Y. Tang, M.-P. Cherrier, S. Choi, J. Salvino and R. Labaudiniere, Tetrahedron Lett., 2000, 41, 1889–1893.
185. C. Hulme and J. Dietrich, Mol. Diversity, 2009, 13, 195–207.
186. R. R. Weikel, J. P. Hallett, C. L. Liotta and C. A. Eckert, Ind. Eng. Chem. Res., 2007, 46, 5252–5257.
187. A.-H. Liu, L.-N. He, F. Hua, Z.-Z. Yang, C.-B. Huang, B. Yu and B. Li, Adv. Synth. Catal., 2011, 353, 3187–3195.
188. R. Ma, C.-B. Huang, A.-H. Liu, X.-D. Li and L.-N. He, Catal. Sci. Technol., 2014, 4, 4308–4312.
189. H.-W. Lin, C. H. Yen, H. Hsu and C.-S. Tan, RSC Adv., 2013, 3, 17222–17227.
190. Y. Huang, S. R. Schricker, B. M. Culbertson and S. V. Olesik, J. Macromol. Sci., Part A: Pure Appl.Chem., 2002, 39, 27–38.
191. N. Akiya and P. E. Savage, Chem. Rev., 2002, 102(8), 2725–2750.
192. G. Brunner, J. Supercrit. Fluids, 2009, 47(3), 373–381.
193. G. Brunner, J. Supercrit. Fluids, 2009, 47(3), 382–390.
194. G. Brunner, Hydrothermal and Supercritical Water Processes, Elsevier, Amsterdam, 2014.
195. A. Kruse and E. Dinjus, J. Supercrit. Fluids, 2007, 39, 362–380.
196. A. Kruse and N. Dahmen, J. Supercrit. Fluids, 2015, 96, 36–45.
197. K. Chandler, B. Eason, C. L. Liotta and C. A. Eckert, Ind. Eng. Chem. Res., 1998, 37, 3515–3518.
198. K. Minami, T. Suzuki, T. Aizawa, K. Sue, K. Arai and R. L. Smith, Fluid Phase Equilib., 2007, 257, 177–182.
199. T. S. Light, Anal. Chem., 1984, 56, 1138–1142.
200. P. A. Alemán, C. Boix and M. Poliakoff, Green Chem., 1999, 1, 65–68.
201. S. E. Hunter and P. E. Savage, Ind. Eng. Chem. Res., 2003, 42(2), 290–294.
202. S. E. Hunter and P. E. Savage, Chem. Eng. Sci., 2004, 59(22–23), 4903–4909.
203. S. E. Hunter and P. E. Savage, AIChE J., 2008, 54(2), 516–528.
204. T. Miyazawa and T. Funazukuri, Biotechnol. Prog., 2005, 21(6), 1782–1785.
205. R. L. Orozco, M. D. Redwood, G. A. Leeke, A. Bahari, R. C. D. Santos and L. E. Macaskie, Int. J. Hydrogen Energy, 2012, 37, 6545–6553.
206. S. Kumar, J. Kumar, S. Kaul and S. L. Jain, Ind. Eng. Chem. Res., 2014, 53, 15571–15575.
207. S. Wu, H. Fan, Y. Xie, Y. Cheng, Q. Wang, Z. Zhang and B. Han, Green Chem., 2010, 12, 1215.
208. A. R. C. Morais, A. M. da Costa Lopes and R. Bogel-Łukasik, Chem. Rev., 2015, 115, 3–27.
209. H. Fan, T. Wu, G. Yang, P. Zhang, G. Ding, W. Wang, T. Jiang and B. Han, Catal. Commun., 2013, 33, 42–46.
210. H. Fan, Q. Wang, J. Guo, T. Jiang, Z. Zhang, G. Yang and B. Han, Green Chem., 2012, 14, 1854–1858.
211. J. Shang, W. Sun, L. Zhao and W.-K. Yuan, Chem. Eng. Sci., 2015, 127, 52–59.
212. S. E. Hunter, C. E. Ehrenberger and P. E. Savage, J. Org. Chem., 2006, 71(16), 6229–6239.
213. Y. Nagai, N. Matubayasi and M. Nakahara, Bull. Chem. Soc. Jpn., 2004, 77, 691–697.
214. A. Yamaguchi, N. Hiyoshi, O. Sato, K. K. Bando and M. Shirai, Green Chem., 2009, 11, 48–52.
215. A. Yamaguchi, N. Hiyoshi, O. Sato and M. Shirai, Top. Catal., 2010, 53, 487–491.
216. A. Yamaguchi, N. Hiyoshi, O. Sato and M. Shirai, ACS Catal., 2011, 1(1), 67–69.
217. A. Yamaguchi, N. Hiyoshi, O. Sato and M. Shirai, Catal. Today, 2012, 185(1), 302–305.
218. A. Yamaguchi, N. Hiyoshi, O. Sato, C. V. Rode and M. Shirai, Chem. Lett., 2008, 37(9), 926–927.
219. O. Sato, A. Yamaguchi and M. Shirai, Catal. Commun., 2015, 68, 6–10.
220. A. P. Abbott, E. G. Hope, R. Mistry and A. M. Stuart, Green Chem., 2009, 11, 1536–1540.
221. C. Tortosa-Estorach, N. Ruiz and A. M. Masdeu-Bultó, Chem. Commun., 2006, 2789–2791.
222. O. A. Kizas, D. Y. Antonov, Y. E. Vopilov, I. A. Godovikov, A. S. Peregudov, N. D. Kagramanov, L. N. Bulatnikova, L. N. Nikitin and A. R. Khokhlov, J. Supercrit. Fluids, 2013, 76, 61–66.
223. H. He, C. Qi, X. Hu, Y. Guan and H. Jiang, Green Chem., 2014, 16, 3729–3733.

---