Why is dissolving potassium nitrate in water an endothermic process ?

Dissolving potassium nitrate in water is an endothermic process because the hydration of the ions when the crystal dissolves does not provide as much energy as is needed to break up the lattice.

The enthalpy change of solution or #DeltaH_(sol)# is the enthalpy change when 1 mole of a solute dissolves to form an "infinitely" dilute solution and can be measured experimentally. In this case it refers to:

#KNO_(3(s))+(aq)rarrKNO_(3(aq)#

We measure the strength of the ionic attractions in a lattice by the lattice enthalpy #DeltaH_(L#. This is the enthalpy change when 1 mole of a solid is formed by the coming together of the ions from infinity in the gaseous state. In this case it is the enthalpy change for:

#K_((g))^++NO_(3(g))^(-)rarrKNO_(3(s)#

Forming an ionic lattice from gaseous ions like this is always an exothermic process since bonds are being formed. So #DeltaH_L# is always negative.

When gaseous ions are hydrated they become surrounded by water molecules, so bonds are formed and this is also exothermic.

This is termed enthalpy of hydration or #DeltaH_(hyd)#

So in this case it is the enthalpy change for:

#K_((g))^++(aq)rarrK_((aq))^+#

and:

#NO_(3(g))^(-)+(aq)rarrNO_(3(aq))^-#

We can think of the enthalpy of solution as the sum of the enthalpies of the following two steps:

1. The energy that is put in which will be #-DeltaH_(L)#

2. The energy that is given back which will be the sum of the hydration enthalpies of the cation and anion i.e #DeltaH_(hyd)(K^+)# +#DeltaH_(hyd)(NO_(3)^-)#

I don't have any values to give you but in this example less energy is released than is put so #DeltaH_(sol)# is positive. I suspect this is due to the low charge density on the nitrate (v) ion which results in a low hydration enthalpy.