Laws of thermodynamics
The four laws of thermodynamics define fundamental physical quantities (temperature, energy, and entropy) that characterize thermodynamic systems. The laws describe how these quantities behave under various circumstances, and forbid certain phenomena (such as perpetual motion).
- Zeroth law of thermodynamics: If two systems are both in thermal equilibrium with a third system then they are in thermal equilibrium with each other. This law helps to define the notion of temperature.
- First law of thermodynamics: Heat and work are forms of energy transfer. Energy is invariably conserved, however the internal energy of a closed system may change as heat is transferred into or out of the system or work is done on or by the system. In real systems work does not always leave the system. For example, changes in molecular energy (potential energy), are generally considered to remain within the system. Similarly, the rotational and vibrational energies of polyatomic molecules remain within the system.
- From the above, all the energy associated with a system must be accounted for as heat, work, chemical energy etc., thus perpetual motion machines of the first kind, which would do work without using the energy resources of a system, are impossible.
- Second law of thermodynamics: An isolated system, if not already in its state of thermodynamic equilibrium, spontaneously evolves towards it. Thermodynamic equilibrium has the greatest entropy amongst the states accessible to the system. Perpetual motion machines of the second kind are thus impossible.
- Third law of thermodynamics: 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.
Classical thermodynamics describes the exchange of work and heat between closed systems. It has a special interest in systems that are individually in states of thermodynamic equilibrium. Thermodynamic equilibrium is a condition of systems which are adequately described by only macroscopic variables. Every physical system, however, when microscopically examined, shows apparently random microscopic statistical fluctuations in its thermodynamic variables of state (entropy, temperature, pressure, etc.). These microscopic fluctuations are negligible for systems which are nearly in thermodynamic equilibrium and which are only macroscopically examined. They become important, however, for systems which are nearly in thermodynamic equilibrium when they are microscopically examined, and, exceptionally, for macroscopically examined systems that are in critical states, and for macroscopically examined systems that are far from thermodynamic equilibrium.
Zeroth law 
The zeroth law of thermodynamics may be stated in the following form:
If two systems are both in thermal equilibrium with a third then they are in thermal equilibrium with each other.
The law is intended to allow the existence of an empirical parameter, the temperature, as a property of a system such that systems in thermal equilibrium with each other have the same temperature. The law as stated here is compatible with the use of a particular physical body, for example a mass of gas, to match temperatures of other bodies, but does not justify regarding temperature as a quantity that can be measured on a scale of real numbers.
Though this version of the law is one of the more commonly stated, it is only one of a diversity of statements that are labeled as "the zeroth law" by competent writers. Some statements go further so as to supply the important physical fact that temperature is one-dimensional, that one can conceptually arrange bodies in real number sequence from colder to hotter. Perhaps there exists no unique "best possible statement" of the "zeroth law", because there is in the literature a range of formulations of the principles of thermodynamics, each of which call for their respectively appropriate versions of the law.
Although these concepts of temperature and of thermal equilibrium are fundamental to thermodynamics and were clearly stated in the nineteenth century, the desire to explicitly number the above law was not widely felt until Fowler and Guggenheim did so in the 1930s, long after the first, second, and third law were already widely understood and recognized. Hence it was numbered the zeroth law. The importance of the law as a foundation to the earlier laws is that it allows the definition of temperature in a non-circular way without reference to entropy, its conjugate variable. Such a temperature definition is said to be 'empirical'.
First law 
The first law of thermodynamics may be stated thus:
- For any process, the increase in internal energy of a body is equal to the heat supplied to the body minus work done by the body.
- For a thermodynamic cycle, the heat supplied to a closed system, minus that removed from it, equals the net work done by the system.
- The change in internal energy of a closed system for any process of interest between an initial and a final state of internal thermodynamic equilibrium is equal to the change in internal energy for a reference process consisting only of adiabatic work that goes from that initial to that final state.
More specifically, the First Law encompasses several principles:
- This states that energy can be neither created nor destroyed. However, energy can change forms, and energy can flow from one place to another. The total energy of an isolated system remains the same.
- The concept of internal energy and its relationship to temperature.
- If a system has a definite temperature, then its total energy has three distinguishable components. If it is in motion, it has kinetic energy. If it is in a field (e.g. gravity), it has potential energy. And it has internal energy which is the sum of the kinetic energy of microscopic motions of its constituent atoms, and of the potential energy of interactions between them. Other things being equal, the internal kinetic energy increases as the system's temperature increases. The concept of internal energy is the characteristic distinguishing feature of the first law of thermodynamics.
- The flow of heat is a form of energy transfer.
- In other words, a quantity of heat that flows from a hot body to a cold one can be expressed as an amount of energy being transferred from the hot body to the cold one.
- Performing work is a form of energy transfer.
- For example, when a machine lifts a heavy object upwards, some energy is transferred from the machine to the object. The object acquires its energy in the form of gravitational potential energy in this example.
Combining these principles leads to one traditional statement of the first law of thermodynamics: it is not possible to construct a perpetual motion machine which will continuously do work without consuming energy.
Second law 
When two initially isolated systems in separate but nearby regions of space, each in thermodynamic equilibrium in itself but not necessarily with each other, are then allowed to interact, they will eventually reach a mutual thermodynamic equilibrium. The sum of the entropies of the initially isolated systems is less than or equal to the total entropy of the final combination.
This statement of the law recognizes that in classical thermodynamics, the entropy of a system is defined only when it has reached its own internal thermodynamic equilibrium. There are other statements of the law.
The second law refers to a wide variety of processes, reversible and irreversible. Its main import is to tell about irreversibility. In an irreversible process of transfer of matter and energy between two systems, the sum of the entropies of the two systems is greater finally than initially. In the trivial case in which the two systems have equal intensive variables, the sum of the entropies does not change. Apart from this trivial case, all natural processes are irreversible, though reversible processes are a convenient theoretical fiction.
The prime example of irreversibility is in the transfer of heat by conduction or radiation. It was known long before the discovery of the notion of entropy that when two bodies initially of different temperatures come into thermal connection, then heat always flows from the hotter body to the colder one.
The second law tells also about kinds of irreversibility other than heat transfer, for example that of chemical reactions. The notion of entropy is needed to provide that wider scope of the law.
According to the second law of thermodynamics, in a theoretical and fictional reversible heat transfer, an element of heat transferred, δQ, is the product of the temperature (T), both of the system and of the sources or destination of the heat, with the increment (dS) of the system's conjugate variable, its entropy (S)
Entropy may also be viewed as a measure of the lack of physical information about the microscopic details of the motion and configuration of a system. The law asserts that for two given macroscopically specified states of a system, there is a quantity called the difference of entropy between them. This entropy difference defines how much additional microscopic physical information is needed to specify one of the macroscopically specified states, given the macroscopic specification of the other - often a conveniently chosen reference state which may be presupposed to exist rather than explicitly stated. A final condition of a natural process always contains microscopically specifiable effects which are not fully and exactly predictable from the macroscopic specification of the initial condition of the process. This is why entropy increases in natural processes - the increase tells how much extra microscopic information is needed to distinguish the final macroscopically specified state from the initial macroscopically specified state.
Third law 
The third law of thermodynamics is sometimes stated as follows:
At zero temperature the system must be in a state with the minimum thermal energy. This statement holds true if the perfect crystal has only one state with minimum energy. Entropy is related to the number of possible microstates according to S = kBln(Ω), where S is the entropy of the system, kB Boltzmann's constant, and Ω the number of microstates (e.g. possible configurations of atoms). At absolute zero there is only 1 microstate possible (Ω=1) and ln(1) = 0.
A more general form of the third law that applies to systems such as glasses that may have more than one minimum energy state:
The entropy of a system approaches a constant value as the temperature approaches zero.
The constant value (not necessarily zero) is called the residual entropy of the system. combining the above the traditional definition is: at absolute zero temperature the entropy is zero for perfectly crystalline solids.
Count Rumford (born Benjamin Thompson) showed, about 1797, that mechanical action can generate indefinitely large amounts of heat, so challenging the caloric theory. The historically first established thermodynamic principle which eventually became the second law of thermodynamics was formulated by Sadi Carnot during 1824. By 1860, as formalized in the works of those such as Rudolf Clausius and William Thomson, two established principles of thermodynamics had evolved, the first principle and the second principle, later restated as thermodynamic laws. By 1873, for example, thermodynamicist Josiah Willard Gibbs, in his memoir Graphical Methods in the Thermodynamics of Fluids, clearly stated the first two absolute laws of thermodynamics. Some textbooks throughout the 20th century have numbered the laws differently. In some fields removed from chemistry, the second law was considered to deal with the efficiency of heat engines only, whereas what was called the third law dealt with entropy increases. Directly defining zero points for entropy calculations was not considered to be a law. Gradually, this separation was combined into the second law and the modern third law was widely adopted.
See also 
- Conservation law
- Entropy production
- Ginsberg's theorem
- Heat death of the universe
- Laws of science
- Table of thermodynamic equations
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Further reading 
- Atkins, Peter, 2007. Four Laws That Drive the Universe. OUP Oxford.
- Goldstein, Martin, and Inge F., 1993. The Refrigerator and the Universe. Harvard Univ. Press. A gentle introduction.