Solution heat of

The heat of solution of a solid or liquid substance can be determined in practically the same manner as that employed 

Experiment.—Determine the Heat of Solution of Potassium Nitrate. 

A quantity of the salt, about 15 gnns., is finely powdered and placed in a test-tube. The latter is weighed, and then placed in a beaker of water surroimded by the protecting cylinders (Fig. 92, B), and the temperature of the water noted. In the calorimeter are placed about 500 gms. of distilled water, and the apparatus then fitted together with thermometer and stirrer as described on p. 279. The same precautions as before as regards temperature readings are taken. 

Wlien the salt has taken the temperature of the water (say after 10 to 15 minutes), the tube is removed, roughly dried, and the contents emptied into th.e water in the calorimeter. Since the accuracy of the determination depends to a considerable extent on the rapidity with which the solid dissolves, it is of importance that the salt should have been finely powdered, and that the stirring should be fairly vigorous. In this case a motor-driven stirrer is preferable to a hand-stirrer. The weight of salt taken is determined by weighing the tube before and after the addition of salt to the water. 

In the above method, it may of course happen that the salt has not the same temperature as the water, so that a 

Defining AH = -Et(ls), the calculated value (GJ) in water, 1.5 eV, compares well with experiment, 1.7 eV. In ammonia, GJ calculate 1.0 eV, which is within the large uncertainty of the experimental value, 1.7 0.7 eV. The agreement may not be much better thanjortner s (1964) earlier calculation via a continuum model, especially if the contribution of H bonds is explicitly considered. 

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The illustrative data presented in Table VII-3 indicate that the total surface energy may amount to a few tenths of a calorie per gram for particles on the order of 1 /xm in size. When the solid interface is destroyed, as by dissolving, the surface energy appears as an extra heat of solution, and with accurate calorimetry it is possible to measure the small difference between the heat of solution of coarse and of finely crystalline material. 

An excellent example of work of this type is given by the investigations of Benson and co-workers [127, 128]. They found, for example, a value of = 276 ergs/cm for sodium chloride. Accurate calorimetry is required since there is only a few calories per mole difference between the heats of solution of coarse and finely divided material. The surface area of the latter may be determined by means of the BET gas adsorption method (see Section XVII-5). 

The excess heat of solution of sample A of finely divided sodium chloride is 18 cal/g, and that of sample B is 12 cal/g. The area is estimated by making a microscopic count of the number of particles in a known weight of sample, and it is found that sample A contains 22 times more particles per gram than does sample B. Are the specific surface energies the same for the two samples If not, calculate their ratio. 

Still another situation is that of a supersaturated or supercooled solution, and straightforward modifications can be made in the preceding equations. Thus in Eq. IX-2, x now denotes the ratio of the actual solute activity to that of the saturated solution. In the case of a nonelectrolyte, x - S/Sq, where S denotes the concentration. Equation IX-13 now contains AH, the molar heat of solution. 

Microcrystals of SrS04 of 30 A diameter have a solubility product at 25°C which is 6.4 times that for large crystals. Calculate the surface tension of the SrS04-H20 interface. Equating surface tension and surface energy, calculate the increase in heat of solution of this SrS04 powder in joules per mole. 

The adsorption of nonelectrolytes at the solid-solution interface may be viewed in terms of two somewhat different physical pictures. In the first, the adsorption is confined to a monolayer next to the surface, with the implication that succeeding layers are virtually normal bulk solution. The picture is similar to that for the chemisorption of gases (see Chapter XVIII) and arises under the assumption that solute-solid interactions decay very rapidly with distance. Unlike the chemisorption of gases, however, the heat of adsorption from solution is usually small it is more comparable with heats of solution than with chemical bond energies. 

The example of Section XI-5B may be completed as follows. It is found that 0 = 0.5 at a butanol concentration of 0.3 g/100 cm. The heat of solution of butanol is 25 cal/g. The molecular area of adsorbed butanol is 40 A. Show that the heat of adsorption of butanol at this concentration is about 50 ergs/cm. ... 

A number of methods have been described in earlier sections whereby the surface free energy or total energy could be estimated. Generally, it was necessary to assume that the surface area was known by some other means conversely, if some estimate of the specific thermodynamic quantity is available, the application may be reversed to give a surface area determination. This is true if the heat of solution of a powder (Section VII-5B), its heat of immersion (Section X-3A), or its solubility increase (Section X-2) are known. 

Procedure. Calculate the heats of solution of the two species, KF and KF HOAc, at each of the four given molalities from a knowledge of the heat capacity. Calculate the enthalpy of solution per mole of solute at each concentration. Find... 

The principles outlined so far may be used to calculate the tower height as long as it is possible to estimate the temperature as a function of Hquid concentration. The classical basis for such an estimate is the assumption that the heat of solution manifests itself entirely in the Hquid stream. It is possible to relate the temperature increase experienced by the Hquid flowing down through the tower to the concentration increase through a simple enthalpy balance, equation 68, and thus correct the equiHbrium line in ajy—a diagram for the heat of solution as shown in Figure 9. 

Genera.1 Ca.se, The simple adiabatic model just discussed often represents an oversimplification, since the real situation implies a multitude of heat effects (/) The heat of solution tends to increase the temperature and thus to reduce the solubihty. 2) In the case of a volatile solvent, partial solvent evaporation absorbs some of the heat. (This effect is particularly important when using water, the cheapest solvent.) (J) Heat is transferred from the hquid to the gas phase and vice versa. (4) Heat is transferred from both phase streams to the shell of the column and from the shell to the outside or to cooling cods. 

Physical Properties. Physical properties of anhydrous hydrogen fluoride are summarized in Table 1. Figure 1 shows the vapor pressure and latent heat of vaporization. The specific gravity of the Hquid decreases almost linearly from 1.1 at —40°C to 0.84 at 80°C (4). The specific heat of anhydrous HF is shown in Figure 2 and the heat of solution in Figure 3. 

Fig. 3. Heat of solution per gram of anhydrous hydrogen fluoride in water when mixed to the final concentration shown in wt % of HF (16—18).

Other studies which have been reported describe unusual chemistry such as HSO F—Nb(S02F) systems (42). Also the unique properties of fluorosulfuric acid have been found to provide unusual solvent systems, which can vary properties such as acidity, heats of solution, enthalpy, and heats of neutralization (43). 

In general, the peilluoioepoxides have boiling points that are quite similar to those of the corresponding fluoroalkenes. They can be distinguished easily from the olefins by it spectroscopy, specifically by the lack of olefinic absorption and the presence of a characteristic band between 1440 and 1550 cm . The nmr spectra of most of the epoxides have been recorded. Litde physical property data concerning these compounds have been pubhshed (Table 1). The stmcture of HFPO by electron diffraction (13) as well as its solubility and heats of solution in some organic solvents have been measured (14,15). 

At ordinary temperatures, formaldehyde gas is readily soluble in water, alcohols, and other polar solvents. Its heat of solution in water and the lower ahphatic alcohols is approximately 63 kJ/mol (15 kcal/mol). The reaction of unhydrated formaldehyde with water is very fast the first-order rate constant... 

Equations 17—20 result from contact between hot metal and slag, and the sulfur and carbon come dissolved in the hot metal. Likewise, the manganese, siUcon, and phosphoms which are produced are dissolved into the hot metal. The heats of solution for these elements in some cases depend on concentration, and are not included in the heats of reaction Hsted above. The ratio of the concentration of the oxide (or element for sulfur) in the slag to the concentration of the element in the hot metal is the partition ratio, and is primarily a function of slag chemistry and temperature. 

AH gg = —43.03 kJ/mol ( — 10.28 kcal/mol) including heat of solution, at standard state m = V) and may require a heat sink to prevent boiling of the reaction mixture. A 30% by weight suspension of MgO in 20°C water boils in the absence of any heat sink. The time to reach boiling is dependent on the reactivity of the MgO raw material, and this time can be only several hours for the more reactive grades of MgO. Investigations of the kinetics of formation of magnesium hydroxide by hydration of MgO have been reported (79). 

The solubihty of the ammonium haUdes in water also increases with increasing formula weight. For ammonium chloride, the integral heat of solution to saturation is 15.7 kj /mol (3.75 kcal/mol) at saturation, the differential heat of solution is 15.2 kj /mol (3.63 kcal/mol). The solubihty of all three salts is given in Table 1 (7). 

Ammonium nitrate has a negative heat of solution in water, and can therefore be used to prepare freezing mixtures. Dissolution of ammonium nitrate in anhydrous ammonia, however, is accompanied by heat evolution. In dilute solution the heat of neutralization of nitric acid using ammonia is 51.8 kj/mol (12.4 kcal/mo). 

The solubihty of ammonium sulfate in 100 g of water is 70.6 g at 0°C and 103.8 g at 100°C. It is insoluble in ethanol and acetone, does not form hydrates, and dehquesces at only about 80% relative humidity. The integral heat of solution of ammonium sulfate to saturation in water is 6.57 kj/mol (1.57... 

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