# If everything tends to entropy, why does water spontaneously evaporate? [closed]

Knowing that the second law of thermodynamics states that everything tends to entropy, why does water and other volatile substances spontaneously evaporate, getting energy to do so from the surface they were touching? Wouldn't that unobey said law, as the substance would get more energetic and the surface it was previously touching less energetic, decreasing the entropy?

P.S.: Im sorry if this question is too basic or of I dont understand something someone says, as Im still in grade ten and not so knowledgeable about chemistry and physics.

• "Everyting tends to entropy" has no scientific meaning. The sentence "tend to entropy" has no meaning. Entropy is a state function. Nothing can "tend to a state function". Commented Mar 1, 2023 at 21:50
• Evaporation or water (or anything else, for that matter) certainly increases entropy. Commented Mar 1, 2023 at 22:05

## 2 Answers

There's actually two factors at play here, which are being confused. The relevant formula is,

$$\Delta G = \Delta H - T\Delta S$$

The total energy of a system is a combination of its entropy (S) and its enthalpy (H). Entropy can be seen as the net disorder, the total amount of possibilities that every component in the system has (position, velocity, etc.). Enthalpy is effectively the energy inherent in all of the chemical bonds that have formed.

As a liquid H2O has a high degree of enthalpy, from all of hydrogen bonds between molecules. This is why evaporation makes the "surface it was previously touching less energetic". Breaking these bonds took energy, cooling the system.

As a gas there are fewer bonds between molecules, so the enthalpy is worse. However, the gas molecules can move in all directions. The more kinetic energy they have the faster they can move and the more possibilities they have, favouring the entropy (the $$T\Delta S$$ in the above equation). Whether H2O is a liquid or gas therefore depends on the temperature. The boiling point (100 Celsius) is the point where H2O as a gas is overall lower in free energy/$$\Delta G$$ than H2O as a liquid.

As an aside, a system in general will always seek to maximize entropy, because that's the statistically most likely outcome. That's what's meant by the second law of thermodynamics.

The second law of thermodynamics actually talks about the total (net) entropy of a system. The water body and the surface that it is in contact with constitutes a system. When water evaporates taking the energy from the contacted body, it's entropy surely increase with a accompanying decrease in the entropy of that surface. But the entropy increase in water is significantly greater than the decrease in the entropy of the surface. This is because water is a liquid and the entropy change it can have is greater due to the characteristic properties of a typical liquid as compared to the change in entropy of the solid surface (solid being already low in entropy).

The result is that the total entropy change of the system of both is positive that is it's entropy has increased.