The first law of thermodynamics states that the energy of an isolated system is conserved. Therefore, to change the energy of a system, energy must be transferred to or from the system. Heat and work are the only two mechanisms by which energy can be transferred to or from a control mass. Work performed on a body is, by definition [1] an energy transfer to the body that is due to a change of the external parameters of the body (such as the volume, magnetization, center of mass position in a gravitational field etc.). Heat is the energy transferred to the body in any other way. This definition of heat applies generally: it does not appeal to any notion of thermal equilibrium.
In case of bodies close to thermal equilibrium where notions such as the temperature can be defined, heat transfer can be related to temperature difference between bodies. It is an irreversible process that leads to the bodies coming closer to mutual thermal equilibrium.
The unit for the amount of energy transferred by heat in the International System of Units SI is the joule (J), though the British Thermal Unit and the calorie are still used in the United States. The unit for the rate of heat transfer is the watt (W = J/s).
Heat transfer is a path function (process quantity), as opposed to a point function (state quantity). Heat flows between systems that are not in thermal equilibrium with each other; it spontaneously flows from the areas of high temperature to areas of low temperature. When two bodies of different temperature come into thermal contact, they will exchange internal energy until their temperatures are equalized; that is, until they reach thermal equilibrium. The adjective hot is used as a relative term to compare the object’s temperature to that of the surroundings (or that of the person using the term). The term heat is used to describe the flow of energy. In the absence of work interactions, the heat that is transferred to an object ends up getting stored in the object in the form of internal energy.
Specific heat is defined as the amount of energy that has to be transferred to or from one unit of mass or mole of a substance to change its temperature by one degree. Specific heat is a property, which means that it depends on the substance under consideration and its state as specified by its properties. Fuels, when burned, are converted to molecules with a lower internal energy. The change in energy is heat. Upon changing from one phase to another, a pure substance releases or absorbs heat without its temperature changing. The amount of heat transfer during a phase change is known as latent heat and depends primarily on the substance and its state.
Notation
The total amount of energy transferred through heat transfer is conventionally abbreviated as Q. The conventional sign convention is that when a body releases heat into its surroundings, Q Q 0 (+). Heat transfer rate, or heat flow per unit time, is denoted by:
.
It is measured in watts. Heat flux is defined as rate of heat transfer per unit cross-sectional area, and is denoted q, resulting in units of watts per square metre, though slightly different notation conventions can be used.
Entropy
In 1856, German physicist Rudolf Clausius defined the second fundamental theorem (the second law of thermodynamics) in the mechanical theory of heat (thermodynamics): "if two transformations which, without necessitating any other permanent change, can mutually replace one another, be called equivalent, then the generations of the quantity of heat Q from work at the temperature T, has the equivalence-value:"[2][3]
In 1865, he came to define this ratio as entropy symbolized by S, such that, for a closed, stationary system:
and thus, by reduction, quantities of heat δQ (an inexact differential) are defined as quantities of TdS (an exact differential):
In other words, the entropy function S facilitates the quantification and measurement of heat flow through a thermodynamic boundary.
Definitions
In modern terms, heat is concisely defined as energy in transit. Scottish physicist James Clerk Maxwell, in his 1871 classic Theory of Heat, was one of the first to enunciate a modern definition of “heat”. In short, Maxwell outlined four stipulations on the definition of heat. One, it is “something which may be transferred from one body to another”, as per the second law of thermodynamics. Two, it can be spoken of as a “measurable quantity”, and thus treated mathematically like other measurable quantities. Three, it “can not be treated as a substance”; for it may be transformed into something which is not a substance, e.g. mechanical work. Lastly, it is “one of the forms of energy”. Similar such modern, succinct definitions of heat are as follows:
- In a thermodynamic sense, heat is never regarded as being stored within a body. Like work, it exists only as energy in transit from one body to another; in thermodynamic terminology, between a system and its surroundings. When energy in the form of heat is added to a system, it is stored not as heat, but as kinetic and potential energy of the atoms and molecules making up the system.[4]
- The noun heat is defined only during the process of energy transfer by conduction or radiation.[5]
- Heat is defined as any spontaneous flow of energy from one object to another, caused by a difference in temperature between the objects.[6]
- Heat may be defined as energy in transit from a high-temperature object to a lower-temperature object.[7]
- Heat as an interaction between two closed systems without exchange of work is a pure heat interaction when the two systems, initially isolated and in a stable equilibrium, are placed in contact. The energy exchanged between the two systems is then called heat.[8]
- Heat is a form of energy possessed by a substance by virtue of the vibrational movement, i.e. kinetic energy, of its molecules or atoms.[9] The kinetic energy and heat may formally be equivalent, but they are not identical.
- Heat is the transfer of energy between substances of different temperatures.
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