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rst law of thermodynamics From Wikipedia, the free encyclopedia Thermodynamics

Branches[show] Laws[hide] Zeroth · First · Second · Third Systems[show] System properties[show] Material properties[show] Equations[show] Potentials[show] History and culture[show] Scientists[show] v · d · e The first law of thermodynamics is an expression of the principle of conservation of work. The law states that energy can be transformed, i.e. changed from one form to another, but cannot be created nor destroyed. It is usually formulated by stating that the change in the internal energy of a system is equal to the amount of heat supplied to the system, minus the amount of work performed by the system on its surroundings. Contents [hide] 1 Original statements 2 Description 3 Evidence for the first law of thermodynamics 3.1 Adiabatic processes 3.2 Isothermal diabatic processes 4 State functional formulation 5 Spatially inhomogeneous systems 6 History 7 See also 8 References 9 Further reading 10 External links [edit]Original statements

The first explicit statement of the first law of thermodynamics, by Rudolf Clausius in 1850, referred to cyclic thermodynamic processes. "In all cases in which work is produced by the agency of heat, a quantity of heat is consumed which is proportional to the work done; and conversely, by the expenditure of an equal quantity of work an equal quantity of heat is produced."[1] Clausius stated the law also in another form, this time referring to the existence of a function of state of the system called the internal energy, and expressing himself in terms of a differential equation for the increments of a thermodynamic process. This equation may be translated into words as follows: In a thermodynamic process, the increment in the internal energy of a system is equal to the difference between the increment of heat accumulated by the system and the increment of work done by it.[2] [edit]Description

The first law of thermodynamics was expressed in two ways by Clausius. One way referred to cyclic processes and the inputs and outputs of the system, but did not refer to increments in the internal state of the system. The other way referred to any incremental change in the internal state of the system, and did not expect the process to be cyclic. A cyclic process is one which can be repeated indefinitely often and still eventually leave the system in its original state. In each repetition of a cyclic process, the work done by the system is proportional to the heat consumed by the system. In a cyclic process in which the system does work on its surroundings, it is necessary that some heat be taken in by the system and some be put out, and the difference is the heat consumed by the system in the process. The constant of proportionality is universal and independent of the system and was measured by Joule in 1845 and 1847. In any incremental process, the change in the internal energy is considered due to a combination of heat added to the system and work done by the system. Taking dU as an infinitesimal (differential) change in internal energy, one writes

where δQ and δW are infinitesimal amounts of heat supplied to the system and work done by the system, respectively. Note that the minus sign in front of δW indicates that a positive amount of work done by the system leads to energy being lost from the system. (An alternate convention is to consider the work performed on the system by its surroundings. This leads to a change in sign of the work. This is the convention adopted by many modern textbooks of physical chemistry, such as those by Peter Atkins and Ira Levine, but many textbooks on physics define work as work done by the system.) When a system expands in a quasistatic process, the work done on the environment is the product of pressure (P) and volume (V) change, i.e. PdV, whereas the work done on the system is -PdV. The change in internal energy of the system is:

Work and heat are expressions of actual physical processes which add or subtract energy, while U is a mathematical abstraction that keeps account of the exchanges of energy that befall the system. Thus the term heat for δQ means that amount of energy added as the result of heating, rather than referring to a particular form of energy. Likewise, work energy for δW means "that amount of energy lost as the result of work". Internal energy is a property of the system whereas work done and heat supplied are not. A significant result of this distinction is that a given internal energy change dU can be achieved by, in principle, many combinations of heat and work. The internal energy of a system is not uniquely defined. It is defined only up to an arbitrary additive constant of integration, which can be adjusted to give arbitrary reference zero levels. This non-uniqueness is in keeping with the abstract mathematical nature of the internal energy. [edit]Evidence for the first law of thermodynamics

The first law of thermodynamics is induced from empirically observed evidence. The original discovery of the law was gradual over a period of perhaps half a century or more, and was mostly in terms of cyclic processes.[3] The following is an account in terms of changes of state through compound processes that are not necessarily cyclic, but are composed of segments of two special kinds, adiabatic and isothermal diabatic. [edit]Adiabatic processes