Laws of thermodynamics
The laws of thermodynamics form an axiomatic basis of thermodynamics. They define fundamental physical quantities, such as temperature, energy, and entropy, to describe thermodynamic systems and describe the transport and conversion of heat and work in thermodynamic processes. The four principles, or laws, of thermodynamics are:      • The zeroth law of thermodynamics allows the assignment of a unique temperature to systems which are in thermal equilibrium with each other. • The first law of thermodynamics mandates conservation of energy and states in particular that the flow of heat is a form of energy transfer. • The second law of thermodynamics states that the entropy of an isolated macroscopic system never decreases, or, equivalently, that perpetual motion machines are impossible. • The third law of thermodynamics concerns the entropy of a perfect crystal at absolute zero temperature, and implies that it is impossible to cool a system to exactly absolute zero. Classical thermodynamics describes the thermal interaction of systems that are individually in a state of thermodynamic equilibrium. Thermal equilibrium is a statistical condition of macroscopic systems, while microscopically all systems undergo random fluctuations. The laws of thermodynamics are strictly valid only in the thermodynamic limit when a macroscopic system may be described practically by an infinite number of microscopic states to satisfy the laws of statistics. Every finite system will exhibit statistical fluctuations in their thermodynamic parameters (entropy, temperature, pressure, etc.), but these are negligible for macroscopic systems, only becoming important for microscopic systems. There have been suggestions of additional laws, but none of them achieve the generality of the accepted laws, and they are not mentioned in standard textbooks.       The laws of thermodynamics have become some of the most important fundamental laws in physics and they are applicable in other natural sciences.
A thermodynamic system is said to be in equilibrium when the thermodynamic parameters describing its state cease to change in time. A system is said to be in thermal equilibrium with another system when it is diathermically connected to the other system, (i.e. the transfer of heat is allowed between the systems) and its thermodynamic parameters cease to change in time. The zeroth law of thermodynamics may be stated as follows: If system A and system B are in thermal equilibrium with system C, then system A is in thermal equilibrium with system B The zeroth law implies that thermal equilibrium, viewed as a binary relation, is a Euclidean relation. If we assume that the binary relationship is also reflexive, then it follows that thermal equilibrium is an equivalence relation. Equivalence relations are also transitive and symmetric. The symmetric relationship allows one to speak of two systems being "in thermal equilibrium with each other", which gives rise to a simpler statement of the zeroth law: If two systems are in thermal equilibrium with a third, they are in thermal equilibrium with each other However, this statement requires the implicit assumption of both symmetry and reflexivity, rather than reflexivity alone. The law is also a statement about measurability. To this effect the law allows the establishment of an empirical parameter, the temperature, as a property of a system such that systems in equilibrium with each other have the same
Laws of thermodynamics temperature. The notion of transitivity permits a system, for example a gas thermometer, to be used as a device to measure the temperature of another system. Although the concept of thermodynamic equilibrium is fundamental to thermodynamics, the need to state it explicitly as a law was not widely perceived until Fowler and Planck stated it 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.
The first law of thermodynamics may be stated as in several ways: Energy can be neither created nor destroyed. It can only change forms. In any process in an isolated system, the total energy remains the same. For a thermodynamic cycle the net heat supplied to the system equals the net work done by the system. The first law of thermodynamics states that energy cannot be created or destroyed; rather, the amount of energy lost in a steady state process cannot be greater than the amount of energy gained. This is the statement of conservation of energy for a thermodynamic system. It refers to the two ways that a closed system transfers energy to and from its surroundings – by the processes of heat and mechanical work. The rate of gain or loss in the stored energy of a system is determined by the rates of these two processes. In open systems, the flow of matter is another energy transfer mechanism, and extra terms must be included in the expression of the first law. The first law clarifies the nature of energy. For processes in an isolated system, it is a conserved quantity that is independent of any particular process path, meaning it is independent of the system history. If an isolated system undergoes a thermodynamic cycle, whether it becomes warmer, cooler, larger, or smaller, then it has the same amount of energy each time it returns to a particular state. Energy is a state function, and infinitesimal changes in the energy are exact differentials with respect to differentials of variables of state. The first law may be expressed by several forms of the fundamental thermodynamic relation: Increase in internal energy of a system = heat supplied to the system + work done on the system
where U is the internal energy, Q is heat and W is work. The definition of the work is also often given in terms of the work performed by a system on its surroundings.
This is a statement of conservation of energy. The net change in internal energy is the energy that flows in as heat minus the energy that flows out as the work that the system performs on it environment. This is also often stated as a definition of the amount of heat of a process: Heat supplied to a system = increase in internal energy of the system + work done by the system The definition of work and its sign is a matter of convention in particular fields of science. In either case, a resulting increase of the internal energy of a system is represented by a positive amount of work. The energy Q (heat) is the product of the temperature (T) and it conjugate variable entropy (S), Q = TdS, and similarly work is the product of pressure (p) with volume (V) change, W = -pdV. The internal energy then may be written as
Laws of thermodynamics
The second law of thermodynamics may be summarized as follows: When two isolated systems in separate but nearby regions of space, each in thermodynamic equilibrium in itself, but not in equilibrium with each other at first, are at some time allowed to interact, breaking the isolation that separates the two systems, and they exchange matter or energy, they will eventually reach a mutual thermodynamic equilibrium. The sum of the entropies of the initial, isolated systems is less than or equal to the entropy of the final exchanging systems. In the process of reaching a new thermodynamic equilibrium, entropy has increased, or at least has not decreased. It follows that the entropy of an isolated macroscopic system never decreases. The second law states that spontaneous natural processes increase entropy overall, or in another formulation that heat can spontaneously flow only from a higher-temperature region to a lower-temperature region, but not the other way around. The second law defines entropy, which may be described as a measure of deficiency of information, or uncertainty, about the microscopic details of the motion and configuration of the system, given only predictable reproducibility of bulk or macroscopic behavior as specified by macroscopic variables. For example, one has less knowledge about the separate fragments of a broken cup than about an intact cup, because when the fragments are separated, one does not know exactly whether they will fit together again, or whether perhaps there is a missing shard. Crystals, the most regularly structured form of matter, with considerable predictability of microscopic configuration, as well as predictability of bulk behavior, have small values of entropy; and gases, which behave predictably in bulk even when their microscopic motions are unknown, have high entropy.
The third law of thermodynamics is usually stated as follows: The entropy of a perfect crystal at absolute zero is exactly equal to zero. This is explained in statistical mechanics by the fact that a perfect crystal has only one possible microstate (microscopic state): The locations and energies of every atom in a crystal are known and fixed. Entropy is related to the number of possible microstates, and with only one microstate, the entropy is exactly zero. The third law is also stated in a form that includes non-crystal systems, such as glasses: As temperature approaches absolute zero, the entropy of a system approaches a minimum. The minimum, not necessarily zero, is called the residual entropy of the system.
In about 1797, Benjamin Thompson showed 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.
Laws of thermodynamics
 Guggenheim, E.A. (1985). Thermodynamics. An Advanced Treatment for Chemists and Physicists, seventh edition, North Holland, Amsterdam, ISBN 0-444-86951-4.  Kittel, C. Kroemer, H. (1980). Thermal Physics, second edition, W.H. Freeman, San Francisco, ISBN 0-7167-1088-9.  Adkins, C.J. (1968). Equilibrium Thermodynamics, McGraw-Hill, London, ISBN 0-07-084057-1.  Kondepudi D. (2008). Introduction to Modern Thermodynamics, Wiley, Chichester, ISBN 978-0-470-01598-8.  Lebon, G., Jou, D., Casas-Vázquez, J. (2008). Understanding Non-equilibrium Thermodynamics. Foundations, Applications, Frontiers, Springer, Berlin, ISBN 978-3-540-74252-4.  Chris Vuille; Serway, Raymond A.; Faughn, Jerry S. (2009). College physics (http:/ / books. google. ca/ books?id=CX0u0mIOZ44C& pg=PT355). Belmont, CA: Brooks/Cole, Cengage Learning. p. 355. ISBN 0-495-38693-6. .  For example: Honig suggests an upper limit to temperature: (http:/ / books. google. com/ books?id=TlGk3meEw5MC& pg=PA425); Jørgensen suggests that ecosystems will take advantage of exergy sources (http:/ / books. google. com/ books?id=mrJBEYow14UC& pg=PA164)  De Groot, S.R., Mazur, P. (1962). Non-equilibrium Thermodynamics, North Holland, Amsterdam.  Glansdorff, P., Prigogine, I. (1971). Thermodynamic Theory of Structure, Stability and Fluctuations, Wiley-Interscience, London, ISBN 0-471-30280-5.
• Goldstein, Martin, and Inge F., 1993. The Refrigerator and the Universe. Harvard Univ. Press. A gentle introduction.
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