Applications of Solubility Data - ACS Publications - American

Research: Science and Education

Applications of Solubility Data Reginald P. T. Tomkins Department of Chemical Engineering, New Jersey Institute of Technology, Newark, NJ 07102; [email protected]

Solubility data have many interesting and important applications in science and technology. For example chemical engineers involved with the design of systems for chemical separation and purification processes rely heavily on solubility data for their calculations. Civil engineers require solubility data to design water treatment facilities to remove hazardous chemicals from supplies of drinking water. Drug solubility is a vital parameter for controlling solubility in the blood stream and various body fluids to facilitate drug delivery systems. When selecting solvent systems for chemical reactions chemists examine both the reactant and product solubilities. The fate of chemical pollutants in the environment is partially explained by a careful analysis of solubility data. Also over the past decade or so several models have been advanced to predict solubility (1–3). Those concerned with solution thermodynamics utilize solubility data to test the applications, predictive ability, and limitations of some of these models. At the undergraduate level students are exposed to solubility in a peripheral way in several laboratory experiments by the need to get materials into solution prior to analysis. Some physical chemistry laboratory courses include an experiment on solubility as a function of temperature (4) or the determination of a solubility product constant, Ksp. Solubility processes are also important for the teaching of analytical chemistry (5). The information discussed in this article could possibly be a useful resource for instructors of introductory, qualitative, quantitative, and environmental chemistry courses. Solubility data are used to create or understand qualitative and quantitative analysis schemes that students use in laboratory classes. The American Chemical Society publication ChemMatters has included an interesting application of solubility for use in teaching at different levels. The example on aquarium chemistry focuses on the importance of dissolved oxygen, salinity, and ammonia, nitrite, and nitrate concentrations (6). A recent publication presents a comprehensive treatment on the experimental determination of solubilities and includes some discussion on applications of solubility data (7). Several other publications have focused on experiments involved with solubility and thermodynamics, the determination of enthalpies of solution from solubility measurements, and an understanding of solubility products (8–12). This article will review some of the applications of solubility data that are important for the practicing engineer and scientist. Gas Solubilities The solubilities of gases in liquids has been studied for over a hundred and fifty years and many measurements are recorded in the literature. This topic is of particular interest to ecologists and environmentalists. Water quality is frequently measured by dissolved oxygen (13, 14) and this is determined by measuring the quantity of oxygen dissolved in natural waters. Dissolved oxygen also plays a role in sewage treatment (15–18).

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One of the earliest applications of gas solubilities pertains to diving. The affliction is known as “bends”. This occurs when a diver goes rapidly from deep water (high pressure) to the surface (lower pressure). The rapid decompression causes air, dissolved in blood and other body fluids, to bubble out of solution. These bubbles impair blood circulation and affect nerve impulses. As helium is much less soluble than nitrogen, to minimize the effects of “bends”, mixtures of helium and oxygen are now used in diving tanks, so that much less gas comes out of solution upon decompression. One important application is in the control and study of pollution. Solubilities of pollutant gases are of special interest at the present time. One method of scrubbing effluent gas streams is to dissolve noxious gases in a suitable solvent. The variation of solubility with temperature as well as the absolute value of the solubility is important in such applications if the solvent is to be regenerated by heating. The understanding of the behavior of pollutant gases in the atmosphere requires a knowledge of solubilities at very low partial pressures in fresh water and in sea water, in the presence of other gases. Many gases have anesthetic properties. These depend, to a large extent, upon solubilities in body fluids. Even xenon has been shown to cause anesthesia (19, 20). The understanding of the behavior of anesthetics in common use and the search for new anesthetics requires accurate information on the solubilities in water and in body fat (21, 22). The solubility of a gas in olive oil is always close to that in animal fat under the same conditions of temperature and pressure. The solubility of a gas in olive oil has been investigated for potential medical interest (23–25). Two new inhalation anesthetics introduced in the 1990s that are in current use are desflurane and sevoflurane (26). Gas solubilities have also found application in the formulation of artificial blood or blood substitutes. For example using the fact that oxygen has a high solubility in perfluorinated hydrocarbons has led to the use of these solvents as blood substitutes (27–29). A number of gas–solvent systems have now been studied to pressures far exceeding barometric pressure. The variation of solubility with pressure of some gases such as ammonia, hydrogen sulfide, and sulfur dioxide have been studied below their critical temperatures over much of the possible concentration range. In some cases mole fraction solubility has ranged from zero to one with a variety of non-linear variations of mole fraction solubility with pressure. A knowledge of the solubility of ammonia at high pressures is important in designing new integrated ammonia and urea processes (30). As the solubility of a particular gas depends upon both temperature and pressure, confusion may arise from the various ways in which such solubilities are reported. There are, in addition, sometimes serious discrepancies between measurements by one group of workers and another group because it may be difficult to establish thermodynamic equilibrium between a gas phase and a liquid phase in an apparatus for measuring gas solubilities (7).

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When apparently reliable gas solubility data are applied to an industrial problem it must be borne in mind that, in the application under consideration, equilibrium may not exist between the gas and liquid phase. A detailed discussion of the thermodynamic equilibrium of the gas and liquid phases is presented in a recent text (31). A basic question of concern to environmental scientists as well as individuals and organizations responsible for environmental management is what happens to pollutants once they are released. One general approach to answering questions about the fate of pollutants is environmental modeling. For this, box models (32–35) are often used in which the environment is thought of as a collection of boxes or compartments (lake, atmosphere, soil, ocean, etc. as appropriate for the environment being modeled). The boxes are then connected by equations that describe the rates of production and destruction of pollutant in each box and the rates of pollutant transfer between boxes. A model can then be “run” by specifying initial quantities of pollutant in each box and calculating pollutant distributions at later times. Many models, both simple and complex, of this sort have been created and used to predict the fate of pollutants. Obviously, the usefulness of these models is strongly dependent on the accuracy of the rate equations that connect the boxes. In assessing the fate of pollutants from marine and aquatic oil spills, one of the important transfer processes that must be considered is between water and the atmosphere. Such evaporation rates may need to be known for many different pollutants at a variety of temperature conditions and are difficult to measure directly. Fortunately, a method has been described for estimating the rate of evaporation of slightly soluble contaminants from water to the atmosphere (36). At any temperature for which the aqueous solubility and vapor pressure of a slightly soluble pollutant are known, the ratio of evaporation rates for water and the pollutant from a saturated solution or dispersion can be calculated. Since considerable solubility data are available for binary systems of hydrocarbons in water, evaporation rates for single hydrocarbons from water can be estimated. Extending this approach to multi-component systems such as petroleum in water is also possible since it has been shown that binary system data can be used to predict multi-component system data with an error of 15% or less (37). Another important application is the removal of toxic, sulfur-containing gases from petroleum products and flue gases from furnaces and metallurgical processes (38). Even at low concentrations the sulfur-containing compounds, hydrogen sulfide and sulfur dioxide, are extremely corrosive to many metals and alloys (39) and extremely toxic to plant and animal life. As elsewhere, in North America many tons per day of these gases are removed by appropriate gas treatment units in an effort to maintain a clean environment. The sources of hydrogen sulfide or other sulfur-containing compounds are the crude oil, natural gas, or tar sand bitumen that are subsequently processed and refined to produce a host of petroleum products. The processes for the removal of hydrogen sulfide are relatively well-established and relatively successful. On the other hand, sulfur dioxide from processes such as the combustion of coal or roasting of sulfurbearing ores is produced in immense quantities (40). Although treatment facilities for flue gases are generally available (41), much more needs to be done to control the incidence of the resulting “acid rain”. A knowledge of the solubilities of sulfur

dioxide in various aqueous solutions of certain salts or organic compounds is essential for the design and operation of flue gas treating equipment. There has been a dearth of information to indicate which solvents or mixed solvent solutions would effectively remove the sulfur dioxide pollutant to an acceptably low concentration on a cost-effective basis (42). Thus a good deal of research effort is still involved with the solubility determinations of sulfur dioxide in various potential solvents. Another area of concern to those involved in waste treatment facilities is that of the possible volatilization of pollutants from biological treatment facilities (43). The rate at which a compound volatilizes from the liquid phase is proportional to the difference between (i) the partial pressure of the compound in the gas phase and (ii) the partial pressure that the compound would have in the gas phase if the state of equilibrium existed between the two phases. The proportionality constant is determined by the mass transfer characteristics of a given installation. Many of the organic priority pollutants are hydrophobic compounds with large activity coefficients. As a result, these compounds will attempt to leave the aqueous phase at the first opportunity. If the vapor pressure is sufficient, the compound will volatilize from an aerated reactor. If the vapor pressure is very low, the compound will attach itself to any available surface (e.g., biomass or reactor walls). Aerobic biological treatment then becomes a race between the rate of biodegradation and the rate of other removal mechanisms. Selective Solvation in Mixed Solvents Selective solvation of the ions of an electrolyte by the components of a binary mixed solvent profoundly affects (44, 45) its solubility. Most often, when a salt is heteroselectively solvated by the constituents of a mixed solvent, its solubility will be higher (46) in mixtures of the two solvents than in the pure solvent components. In many cases, the solubility has been observed to pass through a maximum value at an intermediate composition of the mixed solvent. Several examples of enhanced solubility of salts such as Ag2SO4 in water–acetonitrile, Ag2SO4 in methanol–dimethylsulfoxide, and AgBrO3 in methanol– acetonitrile have been reported. Parker et al. (47) have shown that the solubility of Ag2SO4 in water containing 150 g∙L of 3-hydroxy propionitrile and 140 g∙L H2SO4 at 25 °C is 0.25 M as compared to 0.07 M in pure water; that is, Ag2SO4 is much more soluble in acidic aqueous solutions containing organic nitriles and H2SO4. Further, Gibbs transfer energy measurements reveal that the silver ion is preferentially solvated by the nitrile and SO42− by water and these results do not depend upon acidity up to pH 3. These observations formed the basis (47) for suggesting an improved method of electro refining silver over the conventional method. Conventionally, silver is refined from a solution of silver nitrate of low acidity at high current density in a Thum or Moebius Cell. The silver deposit on the cathode is dendritic, non-adherent, and of moderate purity and the electrolysis involves significant energy consumption. These disadvantages are absent in electrolysis involving the new Ag2SO4–H2SO4 bath containing about 45 g∙L of Ag+ as sulfate (which is below the saturation solubility under these conditions), 140 g∙L of H2SO4, and 150 g∙L of the nitrile.

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Solubility studies of gypsum (CaSO4⋅2H2O) in the presence of different substances such as urea (48) or farmyard manure (49), which are known to enhance its solubility in water, are particularly relevant in the field of agricultural chemistry. The increased solubility of gypsum in urea solutions, which is about five times that in pure water, is attributed to the formation of a complex, CaSO4⋅4CO(NH2)2, whereas the chelating ability of Ca2+ ions with the organic materials in farmyard manure leads to an increase in solubility. There is evidence to show that the addition of gypsum improves (49) the quality of alkaline soils that in turn results in an increase of crop yields. It is noteworthy that urea, which is a nitrogenous fertilizer of the highest nitrogen content (46.7%), is of limited utility because of its hygroscopic nature whereas the gypsum–urea complex is more suitable as a fertilizer because of its weaker hygroscopicity compared to that of urea. Solubility Data and Phase Diagrams in Industrial Processes Solubility data are used in almost all industrial treatments involving separation, extraction, purification, and crystallization. Useful thermodynamic information can also be derived from solubility data (50). A knowledge of solubility and a reasonable model for its behavior often permit the extrapolation of solubility behavior into pressure and temperature ranges where solubility data are not available. The different unit operations used in industry such as the addition of reactants, evaporation and dilution, purification to given specifications can be simulated quantitatively by graphical displacements using solubility data diagrams. From solubility data, the mass balance of each operation can often be deduced. Furthermore, if additional data on enthalpies of reaction or enthalpies of dilution are available, an energy balance and, in some instances, a cost balance can be derived. Flow sheets for processes can be constructed and optimized. Process parameters such as amounts of reactants, types of reactors, fluid flow and heat flow behavior, material recycling, and the like can be calculated (51). In industry, several techniques are closely related to solubility behavior. A knowledge of such behavior and the use of such information by means of phase diagrams can often enhance product recovery. Some of the more common techniques are

• distillation (52, 53)

• fractional crystallization (54, 55)

• zone melting (56)

Others, less usual, are

• supercritical extraction (57–59)

• zone crystallization (50)

• multiphase component partitioning (51)

In some cases, the process itself can be improved or redesigned from the knowledge of solubility behavior. For example, such is true for processes involving

• lyophilization (60–62)

• hydrothermal synthesis (63, 64)

• crystal growth (65–66)

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Applications in the Power Industry For nuclear power stations it is necessary to provide information about the solubility of Ar, Kr, Xe, N2, and H2 in light and heavy water over a wide temperature range (67, 68). This information is necessary for various reasons: (i) Solubility data for Kr and Xe are needed to optimize the working conditions for the early detection of fuel element failure based on degassing samples of the coolant to detect the presence of these gaseous fission products. (ii) Ar solubility is necessary to evaluate possible oxygen in-leakage from air using a mass-balance model (69). (iii) N2 and H2 solubilities are necessary to know precisely that the coolant and moderator are under reducing conditions according to specifications. Such problem areas require a careful survey and subsequent critical evaluation of available literature data. For the steam–water cycle of power plants (70) it is necessary to have data for the solubility of CO2 at different alkaline pHs to know the local concentration of this species in various components of the steam–water cycle. A critical literature survey shows data are fairly thorough. This helps to optimize working conditions and to improve the designs of some auxiliary systems for a new plant. Information is also needed for the solubility of CaSO4 at high temperatures, as it relates to scale formation in boilers and in steam generator tubings (71). The available data show a discrepancy between experimental and recommended values of three orders of magnitude at 573 K. A careful thermodynamic evaluation of the data and the procedures of calculation eliminated this discrepancy and it was possible to recommend a value of the solubility supported both experimentally and theoretically. The values of high temperature solubility were of importance in preparing specifications for feed water purity starting from fresh water available in the plant site, thus minimizing the risk of scale formation. Solubilities of Oxides and Hydroxides of Copper, Silver, Gold, Zinc, Cadmium, and Mercury Solubility data on oxides and hydroxides are of interest to manufacturers of batteries (72, 73). The oxides of silver, zinc, cadmium, and to some extent copper are used in several battery systems. In such systems the oxides are soluble to some extent in the electrolytes. The degree of solubility is an important factor in the behavior and understanding of the electrode processes. The mechanism of the electrode reactions is interpreted in terms of the solubilities of the oxides and hydroxides. Therefore, the solubility data are of use in the design of battery systems and their applications in different environments. A second area where oxide and hydroxide data are of value is in the practice of the winning of metals from their ores (74). Some of the processes that are used involve the leaching of the oxides of the metals and therefore solubility data are of importance in the operation of such processes. The third area has to do with recent concerns about the environment (75). The oxides and hydroxides of many of the metals cited above are components of throw-away items, for example, small batteries used in hearing aids, watches, and so forth. The metals are heavy metals and toxic. As components of waste they can enter the environment by means of solubility processes.



Application in Alkali Metal Chlorate Production

cell voltage 3–4 V

A large quantity of solubility data for alkaline earth metal halates and alkali metal halates has already been published (76, 77). Halate compounds are used widely in pure and applied chemistry. Potassium iodate is used as the titrant for iodometric titrations, and the titration is used to determine the solubilities of alkali, alkaline earth, transition, and rare earth metal iodates. Important practical applications of halates include their use in pyrotechnic production and in the paper pulp industry for the generation of the bleaching agent chlorine dioxide. The electrolysis process for the production of sodium chlorate is due to Krebs and Co. Before the development of coated titanium anodes the cells employed graphite electrodes. With the introduction of activated titanium anodes it was possible to improve the efficiency of the process. The essential basic components in a typical chlorate cell system (78) are illustrated in Figure 1. The principal overall chemical reaction in a chlorate cell is given by

NaCl 3H2O

NaClO3 3H2

(1)

current flow 1–100 kA

á

ź Cl2 gas evolved

H2 gas evolved

anode

cathode solution

NaCl: NaClO3: NaOCl: Na2Cr2O7: temperature: pH :

70–310 g/L 0–650 g/L 1–6 g/L 0.5–7 g/L 30–90 °C 5.5–7.0

Figure 1. Simple chlorate cell (redrawn and used with permission from ref 78).

Equation 1 is an extremely simplistic expression of a series of desirable reactions that occur in a chlorate cell and ignores a number of undesirable reactions. When the purified brine is electrolyzed, chlorine is formed at the anode and caustic soda and hydrogen are formed at the cathode (78–81):

Cathode:

Anode:

2H2O 2e

2 Cl

H2 2OH

(2)

Cl2 2e

(3)

The chlorate cell has no diaphragm and the average pH is close to neutral. Therefore, the following reactions rapidly occur in the vicinity of the anode:

Cl2 H2O HOCl

H Cl HOCl H OCl

(4) (5)

The actual formation of chlorate takes place outside the cell according to the overall equation (77–80) (6)

Figure 2. Isothermal diagram of solubility for ternary NaClO3–NaCl– H2O system (used with permission from ref 84).

During electrolysis, part of the sodium chloride contained in the solution is transformed into chlorate. After concentration of the solution by evaporation and cooling, chlorate crystals precipitate and can be separated with a centrifuge. The mother liquor, which still contains a quantity of chlorate, is recycled into the electrolysis process, after resaturation with sodium chloride. The isothermal data of solubilities of the ternary NaClO3–NaCl–H2O system at various temperatures are needed to analyze a good operating condition for the brine treatment, electrolysis, and chlorate evaporation and crystallization. Data are shown in Figure 2, which gives three isotherms at 0, 15, and 40 °C. The changes of the sodium chlorate and chloride concentration in the process of electrolysis, evaporation, crystallization, and mixing with fresh brine at the good operating conditions are indicated by the arrow lines 0–1, 1–2, 2–3 and 3–0, re-

spectively. Information from this figure is applied to check the sodium chlorate and chloride concentrations while operating the chlorate production plant, and the results obtained are used to judge whether the plant is continuously operated at the proper conditions or not. If the concentrations of chemical species in the cell solution are varied from the appropriate conditions, the undesirable reactions will occur at the anode and the cathode. All of the reactions at the anode can be minimized and eliminated by the adequate brine treatment and the appropriate cell design and operating conditions. The cathode also has its share of undesirable reactions resulting in a loss of current efficiency. The

2HClO ClO

ClO 3 2H 2Cl


313


reactions at the cathode are almost completely inhibited by the addition of sodium dichromate to the cell solution. Dichromate also acts as an excellent buffer, maintaining the pH of the cell at an optimum. The isothermal solubility studies of the ternary NaClO3– NaCl–H2O system have been reported by Winteler (82), Billiter (83), Nallet and Paris (84), Oey and Koopman (85), and Arkhipov, Kashina, and Kuzina (86). The isothermal diagram reported by Nallet and Paris (84) is reproduced here (Figure 3), and the composition of saturated solutions at invariant points with the nature of the solid phases is given in Table 1. The results over a wide temperature range between ‒35 °C and 100 °C are shown in the figure. Applications in Gas Scrubbing Solubility of Formaldehyde in Aqueous Ethylene Glycol Mixtures The process involved is the gas phase oxidation of 1,2ethane diol over a heterogeneous catalyst to produce glyoxal, which is used for crease proofing cotton. (CH2OH)2

air Cu/Ag/P

Figure 3. Isothermal diagram of solubility for ternary NaClO3–NaCl– H2O system (used with permission from ref 84).

(CHO)2 CO2 CO

scrubbing the gas stream in a counter-current flow of cold water in a column packed with glass helices (87). The main impurities in the gaseous stream are formaldehyde and unreacted ethylene glycol. Very little glyoxal reaches this stage because it forms strong hydrates with water and is removed in the condensation stage. To prevent build up of formaldehyde in the scrubber solution, it is necessary to “bleed-off ” a continuous but small percentage of the liquor. It is therefore necessary to determine the solubility of formaldehyde in a range of concentrations of ethylene glycol in water to decide on the optimum conditions for efficient removal of formaldehyde from the gaseous stream.

other HCHO organic acids

On the pilot plant scale, the gaseous effluent stream (after passing through a heat exchanger at (15–20 °C) is mixed with air to replenish the oxygen deficiency and recirculated. Obviously, there has to be some “bleed-off ” of spent gases to maintain a constant volume. Partly because of environmental considerations and partly to prevent the build up of unwanted materials in the recycled gas stream, purification is required. This is achieved by

Table 1. Solubility Data of the Ternary NaClO3–NaCl–H2O System at Invariant Points Composition of Saturated Solutions T/°C

Sodium Chloride/ [g/(100 g H2O)]

Mole Fraction (%)

Sodium Chlorate/ [g/(100 g H2O)]

Mole Fraction (%)

Nature of the Solid Phases

‒26.25

23.9

8.89

31.3

6.03

ice + NaCl·2H2O + NaClO3

‒19.2

25.2

9.48

33.1

6.43

NaCl·2H2O + NaClO3

2.40

56.8

9.24

ice + NaClO3

‒19.2

5.45

‒9.8

27.0

10.4

36.0

7.08

NaCl·2H2O + NaClO3

‒5.7

27.7

10.7

37.3

7.37

NaCl·2H2O + NaCl + NaClO3

+10

24.9

10.5

49.9

9.62

NaCl + NaClO3

+30

21.25

10.10

70.6

13.1

NaCl + NaClO3

+50

17.85

9.552

95.8

16.0

NaCl + NaClO3

+70

14.95

8.899

123.8

20.8

NaCl + NaClO3

+100

12.45

8.884

185

28.6

NaCl + NaClO3

Note: The data are from ref 84.

314



Solubility of Bromine in Aqueous Hydrogen Bromide Solutions In this case the particular problem is the recovery of hydrogen bromide and bromine from certain aromatic bromination processes used in the production of flame retardants for plastics, for example, ArH Br2

ArBr HBr

or more specifically, CH 3 HO

OH á 4Br2

C CH 3

Br

Br CH 3

HO

OH á

C Br

CH 3

4HBr

Br

However, on the plant scale quite a large quantity of unreacted bromine is carried over in the gaseous stream creating a problem of effluent collection. Hydrogen bromide is soluble to approximately 60% in water at room temperature but bromine is only slightly soluble. Fortunately, bromine is quite soluble in aqueous hydrobromic acid solutions. From both environmental and economic considerations the total effluent from this plant has to be trapped and recovered. The problem is to obtain a set of data for the solubility of elemental bromine in a range of hydrobromic acid concentrations at different temperatures. Temperature is another variable in this case because the dissolution of hydrogen bromide in water is exothermic and the higher the temperature the scrubbing unit can be operated at while maintaining high efficiency of extraction, the less cooling water is required. Acknowledgments Several people involved with the IUPAC Solubility Data Series have provided the author with particular applications of solubility data over many years. In particular the contributions of L. Clever (Emory University), G. Hefter (Murdoch University), and H. Miyamoto (Niigata University) are acknowledged with pleasure. Literature Cited 1. Liu, H.; Yao, X.; Zhang, R.; Liu, M.; Hu, Z.; Fan, B. J. Phys. Chem. B 2005, 109 (43), 205–265. 2. Portier, S.; Rochelle, C. Chem. Geology 2005, 217 (3–4), 187. 3. Zha, S. K.; Madras, G. Fluid Phase Equil. 2004, 225 (1–2), 59. 4. Daniels, F.; Williams, J. W.; Bender, P.; Alberty, R. A.; Cornwell, C. D.; Harriman, J. E. Experimental Physical Chemistry, 7th ed.; McGraw-Hill Book Co.: New York, 1970.

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