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12.3: Enthalpy of Solution

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Chemistry

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Enthalpy of Solution
 
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12.3: Enthalpy of Solution

There are two criteria that favor, but do not guarantee, the spontaneous formation of a solution:

  1. A decrease in the internal energy of the system (an exothermic change, as discussed in the previous chapter on thermochemistry)
  2. An increased dispersal of matter in the system (which indicates an increase in the entropy of the system, as you will learn about in the later chapter on thermodynamics)

In the process of dissolution, an internal energy change often, but not always, occurs as heat is absorbed or evolved. An increase in matter dispersal always results when a solution forms from the uniform distribution of solute molecules throughout a solvent.

Spontaneous solution formation is favored, but not guaranteed, by exothermic dissolution processes. While many soluble compounds do, indeed, dissolve with the release of heat, some dissolve endothermically. Ammonium nitrate (NH4NO3) is one such example and is used to make instant cold packs for treating injuries. A thin-walled plastic bag of water is sealed inside a larger bag with solid NH4NO3. When the smaller bag is broken, a solution of NH4NO3 forms, absorbing heat from the surroundings (the injured area to which the pack is applied) and providing a cold compress that decreases swelling. Endothermic dissolutions such as this one require greater energy input to separate the solute species than is recovered when the solutes are solvated, but they are spontaneous nonetheless due to the increase in disorder that accompanies the formation of the solution.

This text is adapted from Openstax, Chemistry 2e, Section 11.1: The Dissolution Process.


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Enthalpy Of Solution Exothermic Process Endothermic Process Sodium Hydroxide Ammonium Chloride Enthalpy Change Solution Formation Solute Particles Solvent Particles Attractive Forces Hess's Law

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