Thermodynamics defines macroscopic variables (such as temperature, internal energy, entropy, and pressure) that describe average properties of material bodies and radiation, explains how they are related and by what laws they change with time. Everything that is not a part of the system constitutes the surroundings. The system and surroundings are separated by a boundary Figure 1. If our system is one mole of a gas in a container, then the boundary is simply the inner wall of the container itself. The single property that the boundary must have is that it be clearly defined, so we can clearly say whether a given part of the world is in our system or in the surroundings. If matter is not able to pass across the boundary, then the system is said to be closed; otherwise, it is open. A closed system may still exchange energy with the surroundings unless the system is an isolated one, in which case neither matter nor energy can pass across the boundary.
Thermodynamics makes no distinction between kinetic and potential energy and it does not assume the existence of atoms and molecules. In the context of chemistry, the internal energy is the sum of the kinetic energy of the molecules, and the potential energy represented by the chemical bonds between the atoms and any other intermolecular forces that may be operative.
The first law of thermodynamics, also known as Law of Conservation of Energy, states that energy can be neither be created nor destroyed; it can only be transferred or changed from one form to another. For example: The dissolution of ammonium nitrate in water in a single-use cold pack may appear to destroy energy as the temperature of the cold pack decreases. However, the heat energy is only converted to a different form, chemical energy that is invested in chemical bonds.
A way of expressing the first law of thermodynamics is that any change in the internal energy (∆U) of a system is given by the sum of the heat (q) that flows across its boundaries and the work (w) done on the system by the surroundings:
This law says that there are two kinds of processes, heat and work, that can lead to a change in the internal energy of a system. Since both heat and work can be measured and quantified, this is the same as saying that any change in the energy of a system must result in a corresponding change in the energy of the world outside the system. In other words, energy cannot be created or destroyed. If heat flows into a system or the surroundings to do work on it, the internal energy increases and the sign of q or w is positive. Conversely, heat flow out of the system or work done by the system will be at the expense of the internal energy, and will therefore be negative.
The second law of thermodynamics says that the entropy of any isolated system not in thermal equilibrium almost always increases. Isolated systems spontaneously evolve towards thermal equilibrium—the state of maximum entropy of the system—in a process known as "thermalization". Equivalently, perpetual motion machines of the second kind are impossible. More simply put: the entropy of the world only increases and never decreases.
A simple application of the second law of thermodynamics is that a room, if not cleaned and tidied, will invariably become more messy and disorderly with time - regardless of how careful one is to keep it clean. When the room is cleaned, its entropy decreases, but the effort to clean it has resulted in an increase in entropy outside the room that exceeds the entropy lost.
Third law of thermodynamics states that the entropy of a system approaches a constant value as the temperature approaches zero. The entropy of a system at absolute zero is typically zero, and in all cases is determined only by the number of different ground states it has. Specifically, the entropy of a pure crystalline substance at absolute zero temperature is zero. This statement holds true if the perfect crystal has only one state with minimum energy.