Enthalpy

EntHaLpy heat content (denoted as H or ΔH) is a quotient or description of thermodynamic potential of a system, which can be used to calculate the "useful" work obtainable from a closed thermodynamic system under constant pressure. It is often calculated as a differential sum, describing the changes within exo- and endo thermic reactions, which minimize at equilibrium. 

EntHaLpy H = U + PV has

units of energy, but it doesn’t necessarily have a direct physical interpretation as a quantity of heat it

is just a defined variable that often simplifies the calculations in the solution of practical thermodynamic problems.

EntHaLpy of SoLution The amount of heat that is absorbed or released when a solute is dissolved in a solvent 

ΔHsoln is negative when the solution process releases heat  ΔHsoln is a positive value when the solution absorbs heat  ΔHsoln = 0 in the case where no heat is absorbed or released 

EntHaLpy of SoLution The magnitude of ΔHsoln provides information about relative intermolecular forces of solute, solvent, and solution. When ΔHsoln = 0, the solution is referred to as an ideal solution. 

The

enthalpy of solution is the enthalpy associated with the process when a process enters solution. There can be different types of enthalpy of solution, the most common being:

EntHaLpy of SoLution (a) The integral enthalpy of solution - the enthalpy change when one mole of solute dissolves in a large excess of pure solvent. (b) The differential enthalpy of solution - the enthalpy change when one mole of solute dissolves in a large amount of solution. (c) The enthalpy of hydration is reserved for the special case when the solvent is water.

EntHaLpy of Hydration Enthalpies of hydration are always negative.  They depend on the charge and size of an ion.  They become more negative when ionic charge increases or when ionic radius decreases. When we deal with solvents other than water we use the term enthalpy of solvation, ΔHsolv 

Lattice EntHaLpy The heat evolved when one mole of a crystalline ionic compound forms from its gaseous ions, measured under standard conditions. 

The

magnitude of the lattice enthalpy depends on how closely the ions pack together in the crystal structure. The

more closely they pack the more exothermic the lattice enthalpy. The

higher the charge on the ions and the smaller their size, the more closely they will pack in the crystal structure.

Trends in Lattice EntHaLpy Down group I lattice enthalpies become less exothermic. (The same applies down group II) 

Down group VII the lattice enthalpies become less exothermic, as the heat of formation becomes less exothermic and as partly as the enthalpy of atomization becomes less endothermic, due to the larger size of the atoms, and therefore weaker bonds between them as elements. 

Answers to Questions

How is the heat solution of CaCl2 affected by: a) number of moles of CaCl2? b) number of moles of H2O? Explain the significance.



as mole of calcium chloride increases, heat of solution increases  as moles of water increases, heat of solution also decreases

Plot cation radius against lattice enthalpy for CaCl2, NaCl, and KCl. Explain the observed trend.

The larger the cation radius, the smaller the lattice enthalpy (inversely proportional). 

Explain the significance of the partial molar heats of solution of solute and solvent. The dissolution process involves both bond breaking and bond formation; one may consider the breaking of solute-solvent bonds that contribute to a decrease in energy, and the formation of solute-solvent bonds that contribute to an increase in energy. This energy is computed on a per molar basis, and because the result in dissolution process is a solute-solvent mixture, one must consider the partial molar heats of solution of solute and solvent in calculating the enthalpy of solution. 

Plot anion radius against lattice enthalpy for KCl, Kbr, and KI. Explain the observed trend.

Lattice enthalpy is dependent on the reciprocal of anion radius. 