The Enthalpy of Solution

The enthalpy of solution is quite small for many simple ionic compounds and can be either positive or negative. It is the difference between two large quantities, the sum of the hydration enthalpies and the lattice energy. 

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... 

Read the article on the original research (Emsiey, 1971) and include a commentary on these results in your report for this experiment. Emsiey claims that the enthalpy of solution... 

More fundamental treatments of polymer solubihty go back to the lattice theory developed independentiy and almost simultaneously by Flory (13) and Huggins (14) in 1942. By imagining the solvent molecules and polymer chain segments to be distributed on a lattice, they statistically evaluated the entropy of solution. The enthalpy of solution was characterized by the Flory-Huggins interaction parameter, which is related to solubihty parameters by equation 5. For high molecular weight polymers in monomeric solvents, the Flory-Huggins solubihty criterion is X A 0.5. 

The partial molar entropy of a component may be measured from the temperature dependence of the activity at constant composition the partial molar enthalpy is then determined as a difference between the partial molar Gibbs free energy and the product of temperature and partial molar entropy. As a consequence, entropy and enthalpy data derived from equilibrium measurements generally have much larger errors than do the data for the free energy. Calorimetric techniques should be used whenever possible to measure the enthalpy of solution. Such techniques are relatively easy for liquid metallic solutions, but decidedly difficult for solid solutions. The most accurate data on solid metallic solutions have been obtained by the indirect method of measuring the heats of dissolution of both the alloy and the mechanical mixture of the components into a liquid metal solvent.05... 

We will assume that rAB = (rAA + rBB ), but we will avoid making any simplifying assumption concerning 6 (or eAB ) by calculating it from the experimental excess free energy AGe. The value of 6 so obtained is used to calculate the enthalpy of solution and TAS6. 

Finally, they measured the enthalpy of solution of C HsO in water as a function of concentration and extrapolated to infinite dilution to get a value of -5.84 kJ-mol-1 for the reaction... 

Ais obtained from the enthalpies of solution of HCl(g) in water, extrapolated to infinite dilution. 

The enthalpy of formation of PuCl3(c) is derived from the measurement of the enthalpy of solution of this compound in 6 M HC1 according to reaction (14)... 

The enthalpy of formation of PuB c) rests on three concordant sets of data for the enthalpy of solution of this compound in 02-free 6 M HCl (66), 1 M HC1 and 0.1 M HCl (67) and comparison with the enthalpy of solution of PuCl3(c) in the same media (18). These data yield virtually identical values for the enthalpy of formation of PuBr3(c) and can be averaged as AHf(PuBr3,c) =... 

Potassium nitrate dissolves readily in water, and its enthalpy of solution is +34.9 kj-niol. (a) Does the enthalpy of solution favor the dissolving process (b) Is the entropy change of the system likely to be positive or negative when the salt dissolves (c) Is the entropy change of the system primarily a result of changes in positional disorder or thermal disorder ... 

The change in molar enthalpy when a substance dissolves is called the enthalpy of solution, AF/so. The change can be measured calorimetrically from the heat released or absorbed when the substance dissolves at constant pressure. However,... 

Because the enthalpy of solution is positive, there is a net inflow of energy as heat when the solid dissolves (recall Fig. 8.23b). Sodium chloride therefore dissolves endothermically, but only to the extent of 3 kj-mol-1. As this example shows, the overall change in enthalpy depends on a very delicate balance between the lattice enthalpy and the enthalpy of hydration. 

A hypothetical solution that obeys Raoult s law exactly at all concentrations is called an ideal solution. In an ideal solution, the interactions between solute and solvent molecules are the same as the interactions between solvent molecules in the pure state and between solute molecules in the pure state. Consequently, the solute molecules mingle freely with the solvent molecules. That is, in an ideal solution, the enthalpy of solution is zero. Solutes that form nearly ideal solutions are often similar in composition and structure to the solvent molecules. For instance, methylbenzene (toluene), C6H5CH, forms nearly ideal solutions with benzene, C6H6. Real solutions do not obey Raoult s law at all concentrations but the lower the solute concentration, the more closely they resemble ideal solutions. Raoult s law is another example of a limiting law (Section 4.4), which in this case becomes increasingly valid as the concentration of the solute approaches zero. A solution that does not obey Raoult s law at a particular solute concentration is called a nonideal solution. Real solutions are approximately ideal at solute concentrations below about 0.1 M for nonelectrolyte solutions and 0.01 M for electrolyte solutions. The greater departure from ideality in electrolyte solutions arises from the interactions between ions, which occur over a long distance and hence have a pronounced effect. Unless stated otherwise, we shall assume that all the solutions that we meet are ideal. 

The enthalpy of mixing is the same as the enthalpy of solution, AH ,, but AHmix is used more commonly for the mixing of two liquids. 

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

The enthalpy of solution of ammonium nitrate in water is positive, (a) Does NH4N05 dissolve endothermically or exothermically (b) Write the chemical equation for the dissolving process, (c) Which is larger for NH4NO , the lattice enthalpy or the enthalpy of hydration ... 

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. 

Values for the enthalpy of solution of hydrogen in transition metals at infinite dilution shown in Figure 7.22 are more negative for the early transition metals. It should be noted that the enthalpies of solution in general are functions of the concentration of the solute. Still, the values at infinite dilution are useful when looking for systematic variations, particularly since changes with composition are often limited. 

The basic principle of solution calorimetry is simple. In one experiment the enthalpy of solution of, for example, LaA103(s) [32] is measured in a particular solvent. In order to convert this enthalpy of solution to an enthalpy of formation, a thermodynamic cycle, which gives the formation reaction... 

Figure 10.8 Experimental setup for measurement of the enthalpy of solution at high temperatures.

Energy is needed to break the ionic bonds in the solid salt and energy is liberated forming hydration complexes like VI. We also break some of the natural hydrogen bonds in the water. The overall change in enthalpy is termed the enthalpy of solution, A// olutioni. Typical values are —207 kJmol-1 for nitric acid 34 kJmol-1 for potassium nitrate and —65.5 kJmol-1 for silver chloride. 

The energy necessary to dissolve 1 mol of solute is called the enthalpy of solution Absolution) (cf- P- 125). A value of AH can be estimated by analysing the solubility s of a solute (which is clearly a function of Repartition)) w t 1 temperature T. 

Worked Example 5.5 Calculate the enthalpy of solution A blution) from the following solubilities s of potassium nitrate as a function of temperature T. Values of s were obtained from solubility experiments. 

---