# Endothermic processes

## Key Questions

Because the driving force of chemical change is NOT enthalpy but ENTROPY....

#### Explanation:

I can quote TWO examples of spontaneous endothermic change, and the first has a highly practical application. Have you ever seen those cold packs used in sports-medicine? These consist of a pack of SOLID ammonium nitrate, and there is a little blister of water. To apply, the blister is broken, AND ALL the ammonium nitrate goes up into solution...and this is occasioned with SUBSTANTIAL HEAT LOSS FROM the surroundings...i.e. $\text{spontaneous endothermic change}$. This is topically applied to the sprain, wound, swelling of an injured athlete..

$N {H}_{4} C l \left(s\right) + \Delta \stackrel{{H}_{2} O \left(l\right)}{\rightarrow} N {H}_{4}^{+} + C {l}^{-}$

And a reaction that may be demonstrated in a first year lab is the reaction between barium hydroxide hydrate and ammonium thiocyanate…

$B a {\left(O H\right)}_{2} \cdot 8 {H}_{2} O + 2 N {H}_{4} {\left(S C N\right)}_{2} \left(s\right) + \Delta \rightarrow B a {\left(S C N\right)}_{2} \left(a q\right) + 2 N {H}_{3} \left(g\right) \uparrow + 10 {H}_{2} O \left(l\right)$

And here the release of ammonia gas provides an entropic driving force to the reaction...

For both reactions, a lot of heat is taken in from the surroundings, and you may see the sides of the flask become frosty as water condenses on the surface....

How does entropy increase in the given examples...?