Dissolution enthalpy change

Measuring enthalpy changes for the dissolution of hydrocarbons, such as alkanes, in water shows that heat is evolved, i.e., A/f is negative and energetically water and alkanes attract each other. However, such attraction does not make alkanes soluble in water to any appreciable extent. This is because the free energy change AGsomtion opposes the process and is positive. 

Thus, the enthalpy change is about half that for the dissolution of ammonium nitrate in water. An estimate of the temperature achieved by the ice pack can be calculated knowing the amounts of water and salt present and assuming that the salt dissolves completely. 

You just learned about the enthalpy changes that are associated with phase changes. Another type of physical change that involves a heat transfer is dissolution. When a solute dissolves in a solvent, the enthalpy change that occurs is called the enthalpy of solution, Affsoin- Dissolution can be either endothermic or exothermic. 

K. Solve this problem in two parts. First, calculate the amount of heat released during dissolution of the NaOH by determining the number of moles of NaOH then multiply that by the enthalpy change given for NaOH in the problem (44.2 kJ) ... 

The dissolution of a solute in a solvent has associated with it a free-energy change, AG = AH — TAS. The enthalpy change is the heat of solution (AHsoin), and the entropy change is the entropy of solution (ASsoin). Heats of solution can be either positive or negative, depending on the relative strengths of solvent-solvent, solute-solute, and solvent-solute intermolecular forces. Entropies of solution are usually positive because disorder increases when a pure solute dissolves in a pure solvent. 

The dissolution reaction (change in free enthalpy Ag1 ) implies the breaking of too Si-O-Si bonds ( per silicon atom) and the formation of four silanol groups (Kith disappearance of tno Hater molecules). On the other hand the polymerization reaction (change in free enthalpy AG) corresponds to the formation of a Si-0-Si bond (and a Hater molecule) and to the disappearance of tno silanol groups. It can thus be nritten to a first approximation ... 

In a coffee cup calorimeter, 1.60 g of NH4NO3 is mixed with 75.0 g of water at an initial temperature of 25.00°C. After dissolution of the salt, the final temperature of the calorimeter contents is 23.34°C. Assuming the solution has a heat capacity of 4.18 J °C 1 g-1 and assuming no heat loss to the calorimeter, calculate the enthalpy change for the dissolution of NH4NO3 in units of kj/mol. 

Hildebrand [12] showed that the enthalpy change per unit volume for endothermic dissolution (the usual case for polymers) is expressed as... 

The A//soi is the enthalpy change that occurs on dissolution of one mole of compound in a solvent. The solution microcalorimeter may be used to obtain the enthalpy of solution directly. The change of free energy can be calculated from the concentration using the enthalpy obtained. The change in the entropy of solution AAsoin can then be determined from the Gibbs-Helmholtz equation. " ... 

Above this temperature, the substance retains some of the properties of a liquid, e.g., molecular mobility, and is termed rubbery. Above this temperature, the increase in molecular mobility facilitates spontaneous crystallization into the crystalline form with an exothermic enthalpy change after the glass transition. The use of amorphous forms is attractive, particularly for sparingly soluble compounds because of the enhanced solubility and dissolution rate over the crystalline state leading to increased bioavailability. However, the amorphous state is thermodynamically unstable. The glass transition temperature Tg is lowered by water or other additives, facilitating conversion... 

The dissolution process is favored by a negative enthalpy change but opposed by a decrease in entropy. Decrease in entropy upon dissolution is characteristic of uncharged solutes. The standard free energy change is... 

Enthalpy changes accompany such processes as the dissolution of a solute, the formation of micelles, chemical reaction, adsorption onto solids, vaporisation of a solvent, hydration of a solute, neutralisation of acids and bases, and the melting or freezing of solutes. 

Utilization of the dissolution enthalpies in Eq. (4.36) is justified when the individual phases are adequately diluted in an amount of the chosen solvent. With regard to the error in the solution calorimetry, the enthalpy of mixing and the enthalpy of dissolution could be neglected only when the amount of the solvent in the solution formed has not changed. 

The sum of the two values, the enthalpy of cooling, Acooi f, and the enthalpy of dissolution, AgoiT/, is the so-called relative enthalpy, Hrei, of the sample. From the temperature dependence of the relative enthalpy, the heat capacity as well as all the following enthalpy, changes could be calculated. 

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... 

The enthalpy change on dissolution of H2Se(g) in 0.25 M lithium, sodium, or potassium hydroxide was measured by Fabre [1887FAB] with concordant results. From these measurements carried out at about 287 K the review estimates in Appendix A the enthalpy of formation of HSe to be (14.3 3.2) kJ-mol . This result has not been selected since a proper correction from 287 K to standard conditions could not be made. 

Selivanova and Pakhorukov [61SEL/PAK] measured the integral heat of dissolution of H2Se03(cr), see Appendix A. The enthalpy change of ... 

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