Acetone Enthalpy & Observations I was looking at the enthalpy of various liquids using WolframAlpha and noticed that at 30 °C, while water and ethanol have a positive enthalpy, acetone has a negative enthalpy.

If I'm understanding this correctly, this means that as acetone evaporates, the temperature of the acetone will increase (because the evaporation will release energy, thereby raising the temperature). I started searching up in temperatures to find the zero-crossing, and that appears to be around 56.074 °C.

This suggests that if a sufficient quantity of acetone is left around, it will eventually heat up to 56.074 °C (which interestingly happens to be when its vapor pressure is really close to the STP pressure).

However, if I get acetone on my hand, I get a distinctive "cool" feel — it, like other liquids, appears to be endothermic and cools down. I may be warm-blooded, but I'm not above 56 °C.

Is there something wrong with my understanding? If not, why does the theory seem to contradict my observations?
The absolute values of enthalpy H at a single state point are meaningless. It is only the difference between two different state points that matter. Thus, the value for a single state point can be set to any arbitrary value. Many handbooks set the arbitrary state point so that the values of these properties are positive for most liquid or gas states. 

Apparently, WolframAlpha is setting enthalpy H to zero for the saturated liquid at the normal boiling point (NBP) for acetone but not for water and ethanol.

For ethanol, WolframAlpha is setting specific enthalpy to h = 200 kJ/kg for the saturated liquid at 0 °C (IIR reference state).

For water, the triple point is selected as the reference state, where the internal energy U of saturated liquid is assigned a zero value. Accordingly, the enthalpy H = U + pV of liquid water at the triple point is slightly larger than zero, at a specific enthalpy of about h = 0.6115 J/kg.