Enthalpies of solution for

I.ithium sulfate dissolves exothermically in water, (a) Is the enthalpy of solution for Ei,S04 positive or negative ... 

The importance of the size of the solute relative to that of the solvent mentioned above is evident also from experimental determinations of the extent of solid solubility in complex oxides and from theoretical evaluations of the enthalpy of solution of large ranges of solutes in a given solvent (e.g. a mineral). The enthalpy of solution for mono-, di- and trivalent trace elements in pyrope and similar systems shows an approximately parabolic variation with ionic radius [44], For the pure mineral, the calculated solution energies always show a minimum at a radius close to that of the host cation. 

Enthalpies of Solution for Lanthanide Trichloride Hexahydrates in Aqueous Magnesium Chloride Solutions at 25°C ... 

In many cases the enthalpy of solution for ionic compounds in water is positive. In these cases we find the solution cooling as the solute dissolves. The mixing tendency of entropy is forcing the solution to do work to pull the ions apart, and since in an adiabatic process such work can be done only at the expense of internal energy, the solution cools. If the enthulpy of solution is sufficiently positive, favorable entrapy may not be able to overcome it and the compound will be insoluble. Thus some ionic compounds, such as KCI04, are essentially insoluble in water at room temperature. 

The enthalpy of solution for ammonium nitrate in water is positive, (a) Does NH4N03 dissolve endothermically or exothermically (b) Write the chemical equation for the dissolving process, (c) Which is larger for NH,N03, the lattice enthalpy or the enthalpy of hydration ... 

The second step shown here combines the terms AHZ (for expanding the solvent) and AH3 (for solvent-solute interactions) and is called the enthalpy (heat) of hydration (AH d). This term represents the enthalpy change associated with the dispersal of a gaseous solute in water. Thus the standard enthalpy of solution for dissolving sodium chloride is the sum of AH and AH yd-... 

Solution calorimetry can be used on one level to merely obtain the enthalpy of solution for a given solute, or can be used in a deeper sense to obtain a full thermodynamic description of a system. The determination of solubility data over a defined temperature range can be used to calculate the differential heat of solution of a given polymorphic form. One can subtract the differential heats of solution obtained for the two polymorphs to deduce the heat of transition (A//Trans) between the two forms ... 

DIFFERENTIAL ENTHALPIES OF SOLUTION FOR PHOSGENE IN VARIOUS ORGANIC SOLVENTS AT STANDARD PRESSURE... 

Solution calorimetry can also be used to evaluate amorphous/crystalline content in a binary mixture. The enthalpy of solution for the amorphous compound is an exothermic event, whereas that of the crystalline hydrate is endothermic. Enthalpy of solution is a sum of several thermal events, that is, heat of wetting (incorporating sorption process, such as surface sorption and complexation), disruption of the crystal lattice, and solvation. The order of magnitude of solution enthalpy for the crystalline compound suggests that the disruption of the crystal lattice predominates over the heat of solvation. In addition, the ready solubility of the compotmd in aqueous media is probably governed by entropy considerations. 

A very simple experiment that has been carried out for many electrolytes in water is the measurement of the enthalpy associated with the dissolution of the electrolyte, which is often a solid, in water. This process can be either exothermic or endothermic, and has an enthalpy change which depends on the relative amounts of electrolyte and water. By studying the enthalpy of solution for one mole of electrolyte as a function of the number of moles of water, which increase from one experiment to the next, one can determine the enthalpy of solution associated with the formation of an infinitely dilute solution. In the case of NaCl, the relevant process is... 

TABLE 10.3 Enthalpy of Solution for Common Aerosol Salts at 298 K... 

For RbCl, an enthalpy of solution for infinite aqueous dilution of approximately 16.7 kJ/mol has been repa-ted (A. Sanahuja 1. L. G6mez-Estdvez, Thermochimica Acta, 1989,156, 85) which supports application of the cycle to approximate this value. 

A U. luO-L solution is made by dissolving u.44i gof CaCl2(s) in water, (a) Calculate the osmotic pressure of this solution at 27 "C, assuming that it is completely dissociated into its component ions, (b) The measured osmotic pressure of this solution is 2.56 atm at 27 C. Explain why it is less than the value calculated in (a), and calculate the van t Hoff factor, i, for the solute in this solution. (See the A Closer Look box on Colligative Properties of Electrolyte Solutions in Section 13.5.) (c) The enthalpy of solution for CaCl2 is AH = —81.3 kj/mol. If the final temperature of the solution is 27 °C, what was its initial temperature (Assume that the density of the solution is 1.00 g/mL, that its specific heat is 4.18 J/g-K, and that the solution loses no heat to its surroundings.)... 

Example 5-1. Given that the lattice energy of CsF is —724 kj/mol, use the hydration enthalpies in Table 5.3 to calculate the enthalpy of solution for CsF. 

Infinite Dilution Enthalpies of Solution for Some Ionic Compounds ... 

The stable form of an ionic solid at saturation is often a hydrate. The enthalpy of solution for hydrates is generally positive (endothermic), especially at... 

The difference between the standard energy and enthalpy of solution for process II (see Section 7.2) is quite small. As an example, for argon in water at 25°C, the standard enthalpy of solution is —2000 cal/mole. The value of P A at 1 atm... 

All of the actinides from Th through Cf form dioxides but several of these have not been studied thermodynamically, due in part to their instability and to limited availability (e.g., it is very difficult to prepare multi-milligrams of Cf02 even though such quantities of the isotope are available). Plots of enthalpy of solution for the f elements have been established (Morss 1986) which permit estimating values for the other actinide dioxides. Although binary oxides above the dioxide stoichiometry are known for some of the actinides (Pa, U, Np), little thermodynamic data are available for these oxides. 

Equation 11.4.10 is the desired relation. It shows how a measurement of the molar integral enthalpy change for a solution process that produces solution of a certain molality can be combined with dilution measurements in order to calculate molar integral enthalpies of solution for more dilute solutions. Experimentally, it is sometimes more convenient to carry out the dilution process than the solution process, especially when the pure solute is a gas or solid. 

Figure 15.22 is an enthalpy level diagram that shows the processes which occur when silver chloride, an almost insoluble salt, is placed in water. Figure 15.23 is an energy level diagram that shows the processes which occur when silver fluoride, a very soluble salt, is placed in water. Note the larger positive enthalpy of solution for silver chloride compared to silver fluoride. 

Figure 15.25 shows the energy cycle applied to sodium hydroxide. Dissolving an ionic compound can be considered as two steps step 1 is the breakdown of the lattice to form gaseous ions (the lattice enthalpy) and step 2 is the hydration of these ions, as polar water molecules are attracted to the ions. Step 3 involves the overall enthalpy of solution for sodium hydroxide. 

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