As a student of Thermodynamics, I had a lot of trouble understanding enthalpy. So during my first year of master's, I undertook this task of finding a very simplified understanding of this vast concept.
Some of the common definitions indicated that enthalpy is the energy required to change from one phase to the other. I applied this definition to the first law of thermodynamics: H (Enthalpy) = U (Internal Energy) + PV (work).
Imagine an ice cube taken out of the refrigerator directly put on a heating pan. The process can be described in three simple steps:
1. Initially its temperature remains zero until all of the ice is converted into water (Latent heat of fusion)
2. Next, the temperature of water rises to 100 degC (sensible heat). Here heat provided to the system corresponds to the rise in temperature
3. Temperature remains constant at 100 degC until all the water is converted to vapor (Latent heat of vaporization)
Enthalpy would therefore describe the total heat provided to the system (1+2+3).
The concept of Entropy on the other hand is a little tricky.
Imagine a perfect vacuum cylinder. Suppose a very small amount of gas inserted in one corner of this cylinder. The gas molecules (initially concentrated in one corner) will spread apart occupying the whole cylinder in some time. It is the natural tendency (read entropy) of any physical matter to spread and occupy as much space as possible. Entropy in this case is positive.
Now, if we were to force the molecules in one corner and create vacuum in the rest of the cylinder, we will be required to act against this natural tendency and provide external energy in some form. Entropy for doing this will be negative (?).
Lets take an example of acid + base --> salt + water. When acid comes in contact with a base, the reaction is going to take place naturally. Entropy is positive.
For a reversible reaction, the value of entropy will depend on the direction of the reaction (?).
However, if the entropy is negative for a backward reaction, the reaction simply won't take place.
