Based on only your observations in Part 1, can you say with certainty if the entropy increases or decreases for the dissolution of each salt? Why or why not?
Now look at the results of your calculations for , , and for the dissolution of each salt in Part 1. Are the values you calculated consistent with your observations and data?
Considering both your calculations and observations, for each of the three compounds in Part 1, state whether the dissolution is always spontaneous, spontaneous at low temperatures (enthalpy driven), spontaneous at high temperatures (entropy driven), or never spontaneous.
Salt
Volume of Water (mL)
Mass of Salt (g)
Initial Temperature (°C)
Final Temperature (°C)
Observations
Sodium Chloride
5.0
1.0173
23.1
21.9
completely dissolved, slightly cooler
Potassium Chloride
5.0
1.0233
22.8
15.2
completely dissolved, cold to the touch
Calcium Chloride
5.0
0.9824
23.2
30.7
completely dissolved, noticeably warmer
NaCl - delta H* = 3.98 kJ/mol
delta S* = 0.0432 kJ/mol
delta G = - 8.812 kJ/mol
KCl - delta H* = 18.04 kJ/mol
delta S* = 0.0761 kJ/mol
delta G = - 4.47 kJ/mol
CaCl2 = delta H* = - 82.32 kJ/mol
delta S* = 0.0513 kJ/mol
delta G = - 97.53 kJ/mol