## 4.45

$\Delta S = \frac{q_{rev}}{T}$

Maria Poblete 2C
Posts: 102
Joined: Wed Sep 18, 2019 12:15 am

### 4.45

Potassium nitrate dissolves readily in water, and its enthalpy of solution is +34.9 kJ/mol.
(a) Does the enthalpy of solution favor the dissolving process?

Can someone explain this conceptually? For a, I don't really understand the solution manual's reasoning. It says that because deltaHsystem is positive, deltaSsurr must be negative (because deltaHsystem is negated). I understand this part, but why does the change in enthalpy depend on the entropy of the surroundings? What exactly is the system in this case?

Brian Tangsombatvisit 1C
Posts: 119
Joined: Sat Aug 17, 2019 12:15 am

### Re: 4.45

The system in this case is the potassium nitrate and water. Enthalpy for the system depends on the entropy of the surroundings because the heat going into the system comes from the surroundings (it doesn't just appear out of nowhere). If the surroundings loses heat as a result of heat gain from the system, then the surroundings loses entropy. If the surroundings loses entropy then the process is unfavorable.