Enthalpy of System: Dissolving NaCl in Water

[wintermute++]

Homework Statement

Consider the dissolving of NaCl in water. Assume the system consists of 0.1 mol NaCl and 1 L of water. Considering that the NaCl readily dissolves in the water and that the ions are strongly stabilized by the water molecules, is it safe to conclude that the dissolution of NaCl in water results in a lower enthalpy of the system? Explain your response.

Homework Equations

E = q + w
H = q
possibly
Ek = 0.5mv^2
Epot = 8.99x10^9(Q1)(Q2)/d

The Attempt at a Solution

The change in enthalpy for the dissolution of NaCl is -0.2kJ per mole and for 0.1 mole NaCl is -0.02 kJ. This heat will be transferred to the H2O which is included in the system and so the heat of the system will rise. This conclusion ignores the stabilization effect of the water molecules (something not even covered in the book at this point). I would guess that the kinetic energy of the individual Na+ and Cl- ions requires work from the H2O molecules for stabilization and that the the two values would be equal and thus cancel. So:
E = q --- q = H and q is positive so the enthalpy of the system H = Hfinal - Hinitial, where Hfinal is larger than Hinitial due to the NaCl dissolution.

I really have no idea if what I just wrote makes any sense or not. Please help me.

 

[Zeppos10]

The dissolution of NaCl in water is an endothermic reaction: my data say +4 kJ/mol.
if the system is isolated, then T will drop. If we assume the system is isothermal, heat will enter the system (Q=0.4kJ), and the enthalpy will increase with this amount (dH=Q).
Hence the answer is NO.