The Laws of Thermodynamics

The Three Laws of Thermodynamics

The laws of thermodynamics define fundamental physical quantities (temperature, energy, and entropy) that characterize thermodynamic systems.

System or Surroundings

In order to avoid confusion, scientists discuss thermodynamic values in reference to a system and its surroundings. Everything that is not a part of the system constitutes its surroundings. The system and surroundings are separated by a boundary. For example, if the system is one mole of a gas in a container, then the boundary is simply the inner wall of the container itself. Everything outside of the boundary is considered the surroundings, which would include the container itself.

The boundary must be clearly defined, so one can clearly say whether a given part of the world is in the 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.

The First Law of Thermodynamics

The first law of thermodynamics, also known as Law of Conservation of Energy, states that energy can neither be created nor destroyed; energy can only be transferred or changed from one form to another. For example, turning on a light would seem to produce energy; however, it is electrical energy that is converted.

A way of expressing the first law of thermodynamics is that any change in the internal energy (∆E) 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 surroundings outside the system. In other words, energy cannot be created or destroyed. If heat flows into a system or the surroundings do work on it, the internal energy increases and the sign of q and w are positive. Conversely, heat flow out of the system or work done by the system (on the surroundings) will be at the expense of the internal energy, and q and w will therefore be negative.

The Second Law of Thermodynamics

The second law of thermodynamics says that the entropy of any isolated system always increases. Isolated systems spontaneously evolve towards thermal equilibrium—the state of maximum entropy of the system. More simply put: the entropy of the universe (the ultimate isolated system) only increases and never decreases.

A simple way to think 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.

The Third Law of Thermodynamics

The third law of thermodynamics states that the entropy of a system approaches a constant value as the temperature approaches absolute 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 (perfect order) at absolute zero temperature is zero. This statement holds true if the perfect crystal has only one state with minimum energy.

Spontaneous and Nonspontaneous Processes

Spontaneous processes do not require energy input to proceed, whereas nonspontaneous processes do.

There are two types of processes (or reactions): spontaneous and non-spontaneous. Spontaneous changes, also called natural processes, proceed when left to themselves, and in the absence of any attempt to drive them in reverse. The sign convention of changes in free energy follows the general convention for thermodynamic measurements. This means a release of free energy from the system corresponds to a negative change in free energy, but to a positive change for the surroundings. Examples include:

• a smell diffusing in a room
• ice melting in lukewarm water
• salt dissolving in water
• iron rusting.

The laws of thermodynamics govern the direction of a spontaneous process, ensuring that if a sufficiently large number of individual interactions (like atoms colliding) are involved, then the direction will always be in the direction of increased entropy.

The Second Law of Thermodynamics

The second law of thermodynamics states that for any spontaneous process, the overall ΔS must be greater than or equal to zero; yet, spontaneous chemical reactions can result in a negative change in entropy. This does not contradict the second law, however, since such a reaction must have a sufficiently large negative change in enthalpy (heat energy). The increase in temperature of the reaction surroundings results in a sufficiently large increase in entropy, such that the overall change in entropy is positive. That is, the ΔS of the surroundings increases enough because of the exothermicity of the reaction so that it overcompensates for the negative ΔS of the system. Since the overall ΔS = ΔS_surroundings + ΔS_system, the overall change in entropy is still positive.

Spontaneous Processes

Spontaneity does not imply that the reaction proceeds with great speed. For example, the decay of diamonds into graphite is a spontaneous process that occurs very slowly, taking millions of years. The rate of a reaction is independent of its spontaneity, and instead depends on the chemical kinetics of the reaction. Every reactant in a spontaneous process has a tendency to form the corresponding product. This tendency is related to stability.

Nonspontaneous Processes

An endergonic reaction (also called a nonspontaneous reaction or an unfavorable reaction) is a chemical reaction in which the standard change in free energy is positive, and energy is absorbed. The total amount of energy is a loss (it takes more energy to start the reaction than what is gotten out of it) so the total energy is a negative net result. Endergonic reactions can also be pushed by coupling them to another reaction, which is strongly exergonic, through a shared intermediate.Saul Steinberg from The New Yorker illustrates a nonspontaneous process here.