Chemical Forums
Chemistry Forums for Students => Undergraduate General Chemistry Forum => Topic started by: cvc121 on January 30, 2017, 09:53:41 PM
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Hi,
I am confused about how to determine whether a process is exothermic or endothermic with respect to changes in intermolecular forces. As an example, the dissolution of sodium acetate anhydrous has a literature value of -17.3 kJ/mol (exothermic) for the change in molar enthalpy value. I know that this value is sum of the energy required to overcome lattice energy, an endothermic process, and the energy of forming bonds with water (solute-solvent attractions), an exothermic process. However, within the crystal of sodium acetate anhydrous, there are ionic interactions b/w cations and anions which give rise to a lattice energy. When dissolved, ion-dipole interactions dominate due to interactions between the ions and water molecules. Can anyone clarify as to why these changes in intermolecular interactions result in an exothermic process? I thought that this would be an endothermic process since ionic bonds require lots of energy to break and so the endothermic contribution of lattice energy would surpass the energy released.
Thanks. All help is very much appreciated!
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It's difficult to predict a priori whether it will be exothermic or endothermic. The enthalpy of solution is generally a small difference between large quantities (lattice energy and solvation enthalpy), and relatively small differences in these quantities can lead to relatively large differences in enthalpy of solution. Both lattice and solvation enthalpies typically have magnitudes much greater than 17 kJ/mol.
See this thread: http://www.chemicalforums.com/index.php?topic=81896.msg297960#msg297960
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If you understand the dipole (here of water molecules) as a set of ions (I exaggerate, because their charge is partial, hence less than H+ and O2-), then the solution of an ionic compound in water resembles just any ionic compound, where positive Na is surrounded by negative O and negative acetate by positive H.
Just like in crystalline NaCl, positive Na is surrounded by negative Cl, but just one step farther you meet positive Na again, so you may understand the surroundings of Na as ClNa dipoles too.
So conceptually, the energies need not be very different, and indeed, dissolution can release or absorb heat depending on the compound.
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Thank you for the replies. I now understand that lattice energies and and solvation enthalpies are large quantities and there is a small difference between them. I know that the type of compound determines whether the process will be exothermic or endothermic, but is there any explanation that can account for hydration enthalpy being slighty larger in my example. For instance, does it have to do with the charges of the ions interacting - greater the charges of the ions = greater lattice/hydration energy?