Temperature
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A temperature is a comparative objective measure of hot and cold. It is measured, typically by a thermometer, through the bulk behavior of a thermometric material, detection of heat radiation, or by particle velocity or kinetic energy. It may be calibrated in any of various temperature scales, Celsius, Fahrenheit, Kelvin, etc.
Measurements with a small thermometer, or by detection of heat radiation, can show that the temperature of a body of material can vary from time to time and from place to place within it. If changes happen too fast, or with too small a spacing, within a body, it may be impossible to define its temperature. Thus the concept of temperature in general has an empirical content.
When conditions are such that temperatures are defined, they govern the direction of heat transfer between bodies. Within a body that exchanges no energy or matter with its surroundings, temperature tends to become spatially uniform as time passes. When a path permeable only to heat is open between two bodies, energy always transfers spontaneously as heat from a hotter body to a colder one. The transfer rate depends on the nature of the path. If they are connected by a path permeable only to heat, and no heat flows between them, then the two bodies are equally hot. If changes are slow and spatially smooth enough to allow consistent comparisons of their hotness with other bodies that are respectively in their own states of internal thermodynamic equilibrium, they obey the Zeroth law of thermodynamics and then they have well defined and equal temperatures. Then thermodynamics provides a fundamental physical definition of temperature, on an absolute scale, relying on the second law of thermodynamics.
The kinetic theory offers a valuable but limited account of the behavior of the materials of macroscopic systems. It indicates the absolute temperature as proportional to the average kinetic energy of the random microscopic motions of their constituent microscopic particles such as electrons, atoms, and molecules.
The coldest theoretical temperature is called absolute zero. It can be approached but not reached in any actual physical system. It is denoted by 0 K on the Kelvin scale, −273.15 °C on the Celsius scale, and −459.67 °F on the Fahrenheit scale. In matter at absolute zero, the motions of microscopic constituents are minimal.
Temperature is important in all fields of natural science, including physics, geology, chemistry, atmospheric sciences and biology.
Contents
 1 Effects of temperature
 2 Temperature scales
 3 Thermodynamic approach to temperature
 4 Kinetic theory approach to temperature
 5 Basic theory
 6 Heat capacity
 7 Temperature measurement
 8 Theoretical foundation
 9 Examples of temperature
 10 See also
 11 Notes and references
 12 Further reading
 13 External links
Effects of temperature
Many physical processes are affected by temperature, such as
 physical properties of materials including the phase (solid, liquid, gaseous or plasma), density, solubility, vapor pressure, electrical conductivity
 rate and extent to which chemical reactions occur
 the amount and properties of thermal radiation emitted from the surface of an object
 speed of sound is a function of the square root of the absolute temperature
Temperature scales
{{#invoke:see alsoseealso}} Template:Refimprove section
Temperature scales differ in two ways: the point chosen as zero degrees, and the magnitudes of incremental units or degrees on the scale.
The Celsius scale (°C) is used for common temperature measurements in most of the world. It is an empirical scale. It developed by a historical progress, which led to its zero point Template:Gaps being defined by the freezing point of water, with additional degrees defined so that Template:Gaps was the boiling point of water, both at sealevel atmospheric pressure. Because of the 100 degree interval, it is called a centigrade scale.^{[1]} Since the standardization of the kelvin in the International System of Units, it has subsequently been redefined in terms of the equivalent fixing points on the Kelvin scale, and so that a temperature increment of one degree Celsius is the same as an increment of one kelvin, though they differ by an additive offset of 273.15.
The United States commonly uses the Fahrenheit scale, on which water freezes at 32 °F and boils at 212 °F at sealevel atmospheric pressure.
Many scientific measurements use the kelvin temperature scale (unit symbol K), named in honor of the Scottish physicist who first defined it. It is a thermodynamic or absolute temperature scale. Its zero point, Template:Gaps, is defined to coincide with coldest physicallypossible temperature (called absolute zero). Its degrees are defined through thermodynamics. The temperature of absolute zero occurs at Template:Gaps = Template:Gaps (or −459.67 °F), and the freezing point of water at sealevel atmospheric pressure occurs at Template:Gaps = Template:Gaps.
The International System of Units (SI) defines a scale and unit for the kelvin or thermodynamic temperature by using the reliably reproducible temperature of the triple point of water as a second reference point (the first reference point being 0 K at absolute zero). The triple point is a singular state with its own unique and invariant temperature and pressure, along with, for a fixed mass of water in a vessel of fixed volume, an autonomically and stably selfdetermining partition into three mutually contacting phases, vapour, liquid, and solid, dynamically depending only on the total internal energy of the mass of water. For historical reasons, the triple point temperature of water is fixed at 273.16 units of the measurement increment.
Thermodynamic approach to temperature
Temperature is one of the principal quantities in the study of thermodynamics.
Kinds of temperature scale
There is a variety of kinds of temperature scale. It may be convenient to classify them as empirically and theoretically based. Empirical temperature scales are historically older, while theoretically based scales arose in the middle of the nineteenth century.^{[2]}^{[3]}
Empirically based scales
Empirically based temperature scales rely directly on measurements of simple physical properties of materials. For example, the length of a column of mercury, confined in a glasswalled capillary tube, is dependent largely on temperature, and is the basis of the very useful mercuryinglass thermometer. Such scales are valid only within convenient ranges of temperature. For example, above the boiling point of mercury, a mercuryinglass thermometer is impracticable. Most materials expand with temperature increase, but some materials, such as water, contract with temperature increase over some specific range, and then they are hardly useful as thermometric materials. A material is of no use as a thermometer near one of its phasechange temperatures, for example its boilingpoint.
In spite of these restrictions, most generally used practical thermometers are of the empirically based kind. Especially, it was used for calorimetry, which contributed greatly to the discovery of thermodynamics. Nevertheless, empirical thermometry has serious drawbacks when judged as a basis for theoretical physics. Empirically based thermometers, beyond their base as simple direct measurements of ordinary physical properties of thermometric materials, can be recalibrated, by use of theoretical physical reasoning, and this can extend their range of adequacy.
Theoretically based scales
Theoretically based temperature scales are based directly on theoretical arguments, especially those of thermodynamics, of kinetic theory, and of quantum mechanics. They rely on theoretical properties of idealized devices and materials. They are more or less comparable with practically feasible physical devices and materials. Theoretically based temperature scales are used to provide calibrating standards for practical empirically based thermometers.
The accepted fundamental thermodynamic temperature scale is the Kelvin scale, based on an ideal cyclic process envisaged for a Carnot heat engine.
An ideal material on which a temperature scale can be based is the ideal gas. The pressure exerted by a fixed volume and mass of an ideal gas is directly proportional to its temperature. Some natural gases show so nearly ideal properties over suitable temperature ranges that they can be used for thermometry; this was important during the development of thermodynamics, and is still of practical importance today.^{[4]}^{[5]} The ideal gas thermometer is, however, not theoretically perfect for thermodynamics. This is because the entropy of an ideal gas at its absolute zero of temperature is not a positive semidefinite quantity, which puts the gas in violation of the third law of thermodynamics. The physical reason is that the ideal gas law, exactly read, refers to the limit of infinitely high temperature and zero pressure.^{[6]}^{[7]}^{[8]}
Measurement of the spectrum of electromagnetic radiation from an ideal threedimensional black body can provide an accurate temperature measurement because the frequency of maximum spectral radiance of blackbody radiation is directly proportional to the temperature of the black body; this is known as Wien's displacement law, and has a theoretical explanation in Planck's law and the Bose–Einstein law.
Measurement of the spectrum of noisepower produced by an electrical resistor can also provide an accurate temperature measurement. The resistor has two terminals and is in effect a onedimensional body. The BoseEinstein law for this case indicates that the noisepower is directly proportional to the temperature of the resistor and to the value of its resistance and to the noise bandwidth. In a given frequency band, the noisepower has equal contributions from every frequency, and is called Johnson noise. If the value of the resistance is known then the temperature can be found.^{[9]}^{[10]}
If molecules, or atoms, or electrons, are emitted from a material and their velocities are measured, the spectrum of their velocities often nearly obeys a theoretical law called the Maxwell–Boltzmann distribution, which gives a wellfounded measurement of temperatures for which the law holds.^{[11]} There have not yet been successful experiments of this same kind that directly use the Fermi–Dirac distribution for thermometry, but perhaps that will be achieved in future.^{[12]}
Absolute thermodynamic scale
The Kelvin scale is called absolute for two reasons. One is Kelvin's, that its formal character is independent of the properties of particular materials. The other reason is that its zero is in a sense absolute, in that it indicates absence of microscopic classical motion of the constituent particles of matter, so that they have a limiting specific heat of zero for zero temperature, according to the third law of thermodynamics. Nevertheless, a Kelvin temperature has a definite numerical value, that has been arbitrarily chosen by tradition. This numerical value also depends on the properties of water, which has a gas–liquid–solid triple point that can be reliably reproduced as a standard experimental phenomenon. The choice of this triple point is also arbitrary and by convention. The Kelvin scale is also called the thermodynamic scale.
Definition of the Kelvin scale
The thermodynamic definition of temperature is due to Kelvin.
It is framed in terms of an idealized device called a Carnot engine, imagined to define a continuous cycle of states of its working body. The cycle is imagined to run so slowly that at each point of the cycle the working body is in a state of thermodynamic equilibrium. There are four limbs in such a Carnot cycle. The engine consists of four bodies. The main one is called the working body. Two of them are called heat reservoirs, so large that their respective nondeformation variables are not changed by transfer of energy as heat through a wall permeable only to heat to the working body. The fourth body is able to exchange energy with the working body only through adiabatic work; it may be called the work reservoir. The substances and states of the two heat reservoirs should be chosen so that they are not in thermal equilibrium with one another. This means that they must be at different fixed temperatures, one, labeled here with the number 1, hotter than the other, labeled here with the number 2. This can be tested by connecting the heat reservoirs successively to an auxiliary empirical thermometric body that starts each time at a convenient fixed intermediate temperature. The thermometric body should be composed of a material that has a strictly monotonic relation between its chosen empirical thermometric variable and the amount of adiabatic isochoric work done on it. In order to settle the structure and sense of operation of the Carnot cycle, it is convenient to use such a material also for the working body; because most materials are of this kind, this is hardly a restriction of the generality of this definition. The Carnot cycle is considered to start from an initial condition of the working body that was reached by the completion of a reversible adiabatic compression. From there, the working body is initially connected by a wall permeable only to heat to the heat reservoir number 1, so that during the first limb of the cycle it expands and does work on the work reservoir. The second limb of the cycle sees the working body expand adiabatically and reversibly, with no energy exchanged as heat, but more energy being transferred as work to the work reservoir. The third limb of the cycle sees the working body connected, through a wall permeable only to heat, to the heat reservoir 2, contracting and accepting energy as work from the work reservoir. The cycle is closed by reversible adiabatic compression of the working body, with no energy transferred as heat, but energy being transferred to it as work from the work reservoir.
With this setup, the four limbs of the reversible Carnot cycle are characterized by amounts of energy transferred, as work from the working body to the work reservoir, and as heat from the heat reservoirs to the working body. The amounts of energy transferred as heat from the heat reservoirs are measured through the changes in the nondeformation variable of the working body, with reference to the previously known properties of that body, the amounts of work done on the work reservoir, and the first law of thermodynamics. The amounts of energy transferred as heat respectively from reservoir 1 and from reservoir 2 may then be denoted respectively Q_{1} and Q_{2}. Then the absolute or thermodynamic temperatures, T_{1} and T_{2}, of the reservoirs are defined so that to be such that
Kelvin's original work postulating absolute temperature was published in 1848. It was based on the work of Carnot, before the formulation of the first law of thermodynamics. Kelvin wrote in his 1848 paper that his scale was absolute in the sense that it was defined "independently of the properties of any particular kind of matter." His definitive publication, which sets out the definition just stated, was printed in 1853, a paper read in 1851.^{[13]}^{[14]}^{[15]}^{[16]}
This definition rests on the physical assumption that there are readily available walls permeable only to heat. In his detailed definition of a wall permeable only to heat, Carathéodory includes several ideas. The nondeformation state variable of a closed system is represented as a real number. A state of thermal equilibrium between two closed systems connected by a wall permeable only to heat means that a certain mathematical relation holds between the state variables, including the respective nondeformation variables, of those two systems (that particular mathematical relation is regarded by Buchdahl as a preferred statement of the zeroth law of thermodynamics).^{[17]} Also, referring to thermal contact equilibrium, "whenever each of the systems S_{1} and S_{2} is made to reach equilibrium with a third system S_{3} under identical conditions, the systems S_{1} and S_{2} are in mutual equilibrium."^{[18]} It may viewed as a restatement of the principle stated by Maxwell in the words: "All heat is of the same kind."^{[19]} This physical idea is also expressed by Bailyn as a possible version of the zeroth law of thermodynamics: "All diathermal walls are equivalent."^{[20]} Thus the present definition of thermodynamic temperature rests on the zeroth law of thermodynamics. Explicitly, this present definition of thermodynamic temperature also rests on the first law of thermodynamics, for the determination of amounts of energy transferred as heat.
Implicitly for this definition, the second law of thermodynamics provides information that establishes the virtuous character of the temperature so defined. It provides that any working substance that complies with the requirement stated in this definition will lead to the same ratio of thermodynamic temperatures, which in this sense is universal, or absolute. The second law of thermodynamics also provides that the thermodynamic temperature defined in this way is positive, because this definition requires that the heat reservoirs not be in thermal equilibrium with one another, and the cycle can be imagined to operate only in one sense if net work is to be supplied to the work reservoir.
Numerical details are settled by making one of the heat reservoirs a cell at the triple point of water, which is defined to have an absolute temperature of 273.16 K.^{[21]} The zeroth law of thermodynamics allows this definition to be used to measure the absolute or thermodynamic temperature of an arbitrary body of interest, by making the other heat reservoir have the same temperature as the body of interest.
Temperature an intensive variable
In thermodynamic terms, temperature is an intensive variable because it is equal to a differential coefficient of one extensive variable with respect to another, for a given body. It thus has the dimensions of a ratio of two extensive variables. In thermodynamics, two bodies are often considered as connected by contact with a common wall, which has some specific permeability properties. Such specific permeability can be referred to a specific intensive variable. An example is a diathermic wall that is permeable only to heat; the intensive variable for this case is temperature. When the two bodies have been in contact for a very long time, and have settled to a permanent steady state, the relevant intensive variables are equal in the two bodies; for a diathermal wall, this statement is sometimes called the zeroth law of thermodynamics.^{[22]}^{[23]}^{[24]}
In particular, when the body is described by stating its internal energy U, an extensive variable, as a function of its entropy S, also an extensive variable, and other state variables V, N, with U = U (S, V, N), then the temperature is equal to the partial derivative of the internal energy with respect to the entropy:
Likewise, when the body is described by stating its entropy S as a function of its internal energy U, and other state variables V, N, with S = S (U, V, N), then the reciprocal of the temperature is equal to the partial derivative of the entropy with respect to the internal energy:
The above definition, equation (1), of the absolute temperature is due to Kelvin. It refers to systems closed to transfer of matter, and has special emphasis on directly experimental procedures. A presentation of thermodynamics by Gibbs starts at a more abstract level and deals with systems open to the transfer of matter; in this development of thermodynamics, the equations (2) and (3) above are actually alternative definitions of temperature.^{[27]}
Temperature local when local thermodynamic equilibrium prevails
Real world bodies are often not in thermodynamic equilibrium and not homogeneous. For study by methods of classical irreversible thermodynamics, a body is usually spatially and temporally divided conceptually into 'cells' of small size. If classical thermodynamic equilibrium conditions for matter are fulfilled to good approximation in such a 'cell', then it is homogeneous and a temperature exists for it. If this is so for every 'cell' of the body, then local thermodynamic equilibrium is said to prevail throughout the body.^{[28]}^{[29]}^{[30]}^{[31]}^{[32]}
It makes good sense, for example, to say of the extensive variable U, or of the extensive variable S, that it has a density per unit volume, or a quantity per unit mass of the system, but it makes no sense to speak of density of temperature per unit volume or quantity of temperature per unit mass of the system. On the other hand, it makes no sense to speak of the internal energy at a point, while when local thermodynamic equilibrium prevails, it makes good sense to speak of the temperature at a point. Consequently, temperature can vary from point to point in a medium that is not in global thermodynamic equilibrium, but in which there is local thermodynamic equilibrium.
Thus, when local thermodynamic equilibrium prevails in a body, temperature can be regarded as a spatially varying local property in that body, and this is because temperature is an intensive variable.
Kinetic theory approach to temperature
A more thorough account of this is below at Theoretical foundation.
Kinetic theory provides a microscopic explanation of temperature, based on macroscopic systems' being composed of many microscopic particles, such as molecules and ions of various species, the particles of a species being all alike. It explains macroscopic phenomena through the classical mechanics of the microscopic particles. The equipartition theorem of kinetic theory asserts that each classical degree of freedom of a freely moving particle has an average kinetic energy of k_{B}T/2 where k_{B} denotes Boltzmann's constant. The translational motion of the particle has three degrees of freedom, so that, except at very low temperatures where quantum effects predominate, the average translational kinetic energy of a freely moving particle in a system with temperature T will be 3k_{B}T/2.
It is possible to measure the average kinetic energy of constituent microscopic particles if they are allowed to escape from the bulk of the system. The spectrum of velocities has to be measured, and the average calculated from that. It is not necessarily the case that the particles that escape and are measured have the same velocity distribution as the particles that remain in the bulk of the system, but sometimes a good sample is possible.
Molecules, such as oxygen (O_{2}), have more degrees of freedom than single spherical atoms: they undergo rotational and vibrational motions as well as translations. Heating results in an increase in temperature due to an increase in the average translational kinetic energy of the molecules. Heating will also cause, through equipartitioning, the energy associated with vibrational and rotational modes to increase. Thus a diatomic gas will require more energy input to increase its temperature by a certain amount, i.e. it will have a greater heat capacity than a monatomic gas.
The process of cooling involves removing internal energy from a system. When no more energy can be removed, the system is at absolute zero, though this cannot be achieved experimentally. Absolute zero is the null point of the thermodynamic temperature scale, also called absolute temperature. If it were possible to cool a system to absolute zero, all classical motion of its particles would cease and they would be at complete rest in this classical sense. Microscopically in the description of quantum mechanics, however, matter still has zeropoint energy even at absolute zero, because of the uncertainty principle.
Basic theory
Temperature is a measure of a quality of a state of a material^{[33]} The quality may be regarded as a more abstract entity than any particular temperature scale that measures it, and is called hotness by some writers. The quality of hotness refers to the state of material only in a particular locality, and in general, apart from bodies held in a steady state of thermodynamic equilibrium, hotness varies from place to place. It is not necessarily the case that a material in a particular place is in a state that is steady and nearly homogeneous enough to allow it to have a welldefined hotness or temperature. Hotness may be represented abstractly as a onedimensional manifold. Every valid temperature scale has its own onetoone map into the hotness manifold.^{[34]}^{[35]}
When two systems in thermal contact are at the same temperature no heat transfers between them. When a temperature difference does exist heat flows spontaneously from the warmer system to the colder system until they are in thermal equilibrium. Heat transfer occurs by conduction or by thermal radiation.^{[36]}^{[37]}^{[38]}^{[39]}^{[40]}^{[41]}^{[42]}^{[43]}
Experimental physicists, for example Galileo and Newton,^{[44]} found that there are indefinitely many empirical temperature scales. Nevertheless, the zeroth law of thermodynamics says that they all measure the same quality.
Temperature for bodies in thermodynamic equilibrium
For experimental physics, hotness means that, when comparing any two given bodies in their respective separate thermodynamic equilibria, any two suitably given empirical thermometers with numerical scale readings will agree as to which is the hotter of the two given bodies, or that they have the same temperature.^{[45]} This does not require the two thermometers to have a linear relation between their numerical scale readings, but it does require that the relation between their numerical readings shall be strictly monotonic.^{[46]}^{[47]} A definite sense of greater hotness can be had, independently of calorimetry, of thermodynamics, and of properties of particular materials, from Wien's displacement law of thermal radiation: the temperature of a bath of thermal radiation is proportional, by a universal constant, to the frequency of the maximum of its frequency spectrum; this frequency is always positive, but can have values that tend to zero. Thermal radiation is initially defined for a cavity in thermodynamic equilibrium. These physical facts justify a mathematical statement that hotness exists on an ordered onedimensional manifold. This is a fundamental character of temperature and thermometers for bodies in their own thermodynamic equilibrium.^{[2]}^{[34]}^{[35]}^{[48]}^{[49]}
Except for a system undergoing a firstorder phase change such as the melting of ice, as a closed system receives heat, without change in its volume and without change in external force fields acting on it, its temperature rises. For a system undergoing such a phase change so slowly that departure from thermodynamic equilibrium can be neglected, its temperature remains constant as the system is supplied with latent heat. Conversely, a loss of heat from a closed system, without phase change, without change of volume, and without change in external force fields acting on it, decreases its temperature.^{[50]}
Temperature for bodies in a steady state but not in thermodynamic equilibrium
While for bodies in their own thermodynamic equilibrium states, the notion of temperature requires that all empirical thermometers must agree as to which of two bodies is the hotter or that they are at the same temperature, this requirement is not safe for bodies that are in steady states though not in thermodynamic equilibrium. It can then well be that different empirical thermometers disagree about which is the hotter, and if this is so, then at least one of the bodies does not have a well defined absolute thermodynamic temperature. Nevertheless, any one given body and any one suitable empirical thermometer can still support notions of empirical, nonabsolute, hotness and temperature, for a suitable range of processes. This is a matter for study in nonequilibrium thermodynamics.
Temperature for bodies not in a steady state
When a body is not in a steady state, then the notion of temperature becomes even less safe than for a body in a steady state not in thermodynamic equilibrium. This is also a matter for study in nonequilibrium thermodynamics.
Thermodynamic equilibrium axiomatics
For axiomatic treatment of thermodynamic equilibrium, since the 1930s, it has become customary to refer to a zeroth law of thermodynamics. The customarily stated minimalist version of such a law postulates only that all bodies, which when thermally connected would be in thermal equilibrium, should be said to have the same temperature by definition, but by itself does not establish temperature as a quantity expressed as a real number on a scale. A more physically informative version of such a law views empirical temperature as a chart on a hotness manifold.^{[34]}^{[49]}^{[51]} While the zeroth law permits the definitions of many different empirical scales of temperature, the second law of thermodynamics selects the definition of a single preferred, absolute temperature, unique up to an arbitrary scale factor, whence called the thermodynamic temperature.^{[2]}^{[34]}^{[52]}^{[53]}^{[54]}^{[55]} If internal energy is considered as a function of the volume and entropy of a homogeneous system in thermodynamic equilibrium, thermodynamic absolute temperature appears as the partial derivative of internal energy with respect the entropy at constant volume. Its natural, intrinsic origin or null point is absolute zero at which the entropy of any system is at a minimum. Although this is the lowest absolute temperature described by the model, the third law of thermodynamics postulates that absolute zero cannot be attained by any physical system.
Heat capacity
{{#invoke:see alsoseealso}} When an energy transfer to or from a body is only as heat, state of the body changes. Depending on the surroundings and the walls separating them from the body, various changes are possible in the body. They include chemical reactions, increase of pressure, increase of temperature, and phase change. For each kind of change under specified conditions, the heat capacity is the ratio of the quantity of heat transferred to the magnitude of the change. For example, if the change is an increase in temperature at constant volume, with no phase change and no chemical change, then the temperature of the body rises and its pressure increases. The quantity of heat transferred, ΔQ, divided by the observed temperature change, ΔT, is the body's heat capacity at constant volume, C_{V}.
If heat capacity is measured for a well defined amount of substance, the specific heat is the measure of the heat required to increase the temperature of such a unit quantity by one unit of temperature. For example, to raise the temperature of water by one kelvin (equal to one degree Celsius) requires 4186 joules per kilogram (J/kg)..
Temperature measurement
{{#invoke:see alsoseealso}} Temperature measurement using modern scientific thermometers and temperature scales goes back at least as far as the early 18th century, when Gabriel Fahrenheit adapted a thermometer (switching to mercury) and a scale both developed by Ole Christensen Rømer. Fahrenheit's scale is still in use in the United States for nonscientific applications.
Temperature is measured with thermometers that may be calibrated to a variety of temperature scales. In most of the world (except for Belize, Myanmar, Liberia and the United States), the Celsius scale is used for most temperature measuring purposes. Most scientists measure temperature using the Celsius scale and thermodynamic temperature using the Kelvin scale, which is the Celsius scale offset so that its null point is Template:Gaps = Template:Gaps, or absolute zero. Many engineering fields in the U.S., notably hightech and US federal specifications (civil and military), also use the Kelvin and Celsius scales. Other engineering fields in the U.S. also rely upon the Rankine scale (a shifted Fahrenheit scale) when working in thermodynamicrelated disciplines such as combustion.
Units
The basic unit of temperature in the International System of Units (SI) is the kelvin. It has the symbol K.
For everyday applications, it is often convenient to use the Celsius scale, in which Template:Gaps corresponds very closely to the freezing point of water and Template:Gaps is its boiling point at sea level. Because liquid droplets commonly exist in clouds at subzero temperatures, Template:Gaps is better defined as the melting point of ice. In this scale a temperature difference of 1 degree Celsius is the same as a Template:Gaps increment, but the scale is offset by the temperature at which ice melts (273.15 K).
By international agreement^{[56]} the Kelvin and Celsius scales are defined by two fixing points: absolute zero and the triple point of Vienna Standard Mean Ocean Water, which is water specially prepared with a specified blend of hydrogen and oxygen isotopes. Absolute zero is defined as precisely Template:Gaps and Template:Gaps. It is the temperature at which all classical translational motion of the particles comprising matter ceases and they are at complete rest in the classical model. Quantummechanically, however, zeropoint motion remains and has an associated energy, the zeropoint energy. Matter is in its ground state,^{[57]} and contains no thermal energy. The triple point of water is defined as Template:Gaps and Template:Gaps. This definition serves the following purposes: it fixes the magnitude of the kelvin as being precisely 1 part in 273.16 parts of the difference between absolute zero and the triple point of water; it establishes that one kelvin has precisely the same magnitude as one degree on the Celsius scale; and it establishes the difference between the null points of these scales as being Template:Gaps (Template:Gaps = Template:Gaps and Template:Gaps = Template:Gaps).
In the United States, the Fahrenheit scale is widely used. On this scale the freezing point of water corresponds to 32 °F and the boiling point to 212 °F. The Rankine scale, still used in fields of chemical engineering in the U.S., is an absolute scale based on the Fahrenheit increment.
Conversion
The following table shows the temperature conversion formulas for conversions to and from the Celsius scale. Template:Temperature
Plasma physics
The field of plasma physics deals with phenomena of electromagnetic nature that involve very high temperatures. It is customary to express temperature as energy in units of electronvolts (eV) or kiloelectronvolts (keV). The energy, which has a different dimension from temperature, is then calculated as the product of the Boltzmann constant and temperature, . Then, 1 eV corresponds to Template:Gaps. In the study of QCD matter one routinely encounters temperatures of the order of a few hundred MeV, equivalent to about Template:Gaps.
Theoretical foundation
{{#invoke:see alsoseealso}} Historically, there are several scientific approaches to the explanation of temperature: the classical thermodynamic description based on macroscopic empirical variables that can be measured in a laboratory; the kinetic theory of gases which relates the macroscopic description to the probability distribution of the energy of motion of gas particles; and a microscopic explanation based on statistical physics and quantum mechanics. In addition, rigorous and purely mathematical treatments have provided an axiomatic approach to classical thermodynamics and temperature.^{[58]} Statistical physics provides a deeper understanding by describing the atomic behavior of matter, and derives macroscopic properties from statistical averages of microscopic states, including both classical and quantum states. In the fundamental physical description, using natural units, temperature may be measured directly in units of energy. However, in the practical systems of measurement for science, technology, and commerce, such as the modern metric system of units, the macroscopic and the microscopic descriptions are interrelated by the Boltzmann constant, a proportionality factor that scales temperature to the microscopic mean kinetic energy.
The microscopic description in statistical mechanics is based on a model that analyzes a system into its fundamental particles of matter or into a set of classical or quantummechanical oscillators and considers the system as a statistical ensemble of microstates. As a collection of classical material particles, temperature is a measure of the mean energy of motion, called kinetic energy, of the particles, whether in solids, liquids, gases, or plasmas. The kinetic energy, a concept of classical mechanics, is half the mass of a particle times its speed squared. In this mechanical interpretation of thermal motion, the kinetic energies of material particles may reside in the velocity of the particles of their translational or vibrational motion or in the inertia of their rotational modes. In monatomic perfect gases and, approximately, in most gases, temperature is a measure of the mean particle kinetic energy. It also determines the probability distribution function of the energy. In condensed matter, and particularly in solids, this purely mechanical description is often less useful and the oscillator model provides a better description to account for quantum mechanical phenomena. Temperature determines the statistical occupation of the microstates of the ensemble. The microscopic definition of temperature is only meaningful in the thermodynamic limit, meaning for large ensembles of states or particles, to fulfill the requirements of the statistical model.
In the context of thermodynamics, the kinetic energy is also referred to as thermal energy. The thermal energy may be partitioned into independent components attributed to the degrees of freedom of the particles or to the modes of oscillators in a thermodynamic system. In general, the number of these degrees of freedom that are available for the equipartitioning of energy depend on the temperature, i.e. the energy region of the interactions under consideration. For solids, the thermal energy is associated primarily with the vibrations of its atoms or molecules about their equilibrium position. In an ideal monatomic gas, the kinetic energy is found exclusively in the purely translational motions of the particles. In other systems, vibrational and rotational motions also contribute degrees of freedom.
Kinetic theory of gases
Maxwell and Boltzmann developed a kinetic theory that yields a fundamental understanding of temperature in gases.^{[59]} This theory also explains the ideal gas law and the observed heat capacity of monatomic (or 'noble') gases.^{[60]}^{[61]}^{[62]}
The ideal gas law is based on observed empirical relationships between pressure (p), volume (V), and temperature (T), and was recognized long before the kinetic theory of gases was developed (see Boyle's and Charles's laws). The ideal gas law states:^{[63]}
where n is the number of moles of gas and R = Template:Val is the gas constant.
This relationship gives us our first hint that there is an absolute zero on the temperature scale, because it only holds if the temperature is measured on an absolute scale such as Kelvins. The ideal gas law allows one to measure temperature on this absolute scale using the gas thermometer. The temperature in kelvins can be defined as the pressure in pascals of one mole of gas in a container of one cubic meter, divided by the gas constant.
Although it is not a particularly convenient device, the gas thermometer provides an essential theoretical basis by which all thermometers can be calibrated. As a practical matter it is not possible to use a gas thermometer to measure absolute zero temperature since the gases tend to condense into a liquid long before the temperature reaches zero. It is possible, however, to extrapolate to absolute zero by using the ideal gas law, as shown in the figure.
The kinetic theory assumes that pressure is caused by the force associated with individual atoms striking the walls, and that all energy is translational kinetic energy. Using a sophisticated symmetry argument,^{[64]} Boltzmann deduced what is now called the Maxwell–Boltzmann probability distribution function for the velocity of particles in an ideal gas. From that probability distribution function, the average kinetic energy, E_{k} (per particle), of a monatomic ideal gas is:^{[61]}^{[65]}
where the Boltzmann constant, Template:Var, is the ideal gas constant divided by the Avogadro number, and v_{rms} is the rootmeansquare speed. Thus the ideal gas law states that internal energy is directly proportional to temperature.^{[66]} This direct proportionality between temperature and internal energy is a special case of the equipartition theorem, and holds only in the classical limit of an ideal gas. It does not hold for most substances, although it is true that temperature is a monotonic (nondecreasing) function of internal energy.
Zeroth law of thermodynamics
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When two otherwise isolated bodies are connected together by a rigid physical path impermeable to matter, there is spontaneous transfer of energy as heat from the hotter to the colder of them. Eventually they reach a state of mutual thermal equilibrium, in which heat transfer has ceased, and the bodies' respective state variables have settled to become unchanging.
One statement of the zeroth law of thermodynamics is that if two systems are each in thermal equilibrium with a third system, then they are also in thermal equilibrium with each other.
This statement helps to define temperature but it does not, by itself, complete the definition. An empirical temperature is a numerical scale for the hotness of a thermodynamic system. Such hotness may be defined as existing on a onedimensional manifold, stretching between hot and cold. Sometimes the zeroth law is stated to include the existence of a unique universal hotness manifold, and of numerical scales on it, so as to provide a complete definition of empirical temperature.^{[51]} To be suitable for empirical thermometry, a material must have a monontonic relation between hotness and some easily measured state variable, such as pressure or volume, when all other relevant coordinates are fixed. An exceptionally suitable system is the ideal gas, which can provide a temperature scale that matches the absolute Kelvin scale. The Kelvin scale is defined on the basis of the second law of thermodynamics.
Second law of thermodynamics
{{#invoke:mainmain}} In the previous section certain properties of temperature were expressed by the zeroth law of thermodynamics. It is also possible to define temperature in terms of the second law of thermodynamics which deals with entropy. Entropy is often thought of as a measure of the disorder in a system. The second law states that any process will result in either no change or a net increase in the entropy of the universe. This can be understood in terms of probability.
For example, in a series of coin tosses, a perfectly ordered system would be one in which either every toss comes up heads or every toss comes up tails. This means that for a perfectly ordered set of coin tosses, there is only one set of toss outcomes possible: the set in which 100% of tosses come up the same. On the other hand, there are multiple combinations that can result in disordered or mixed systems, where some fraction are heads and the rest tails. A disordered system can be 90% heads and 10% tails, or it could be 98% heads and 2% tails, et cetera. As the number of coin tosses increases, the number of possible combinations corresponding to imperfectly ordered systems increases. For a very large number of coin tosses, the combinations to ~50% heads and ~50% tails dominates and obtaining an outcome significantly different from 50/50 becomes extremely unlikely. Thus the system naturally progresses to a state of maximum disorder or entropy.
It has been previously stated that temperature governs the transfer of heat between two systems and it was just shown that the universe tends to progress so as to maximize entropy, which is expected of any natural system. Thus, it is expected that there is some relationship between temperature and entropy. To find this relationship, the relationship between heat, work and temperature is first considered. A heat engine is a device for converting thermal energy into mechanical energy, resulting in the performance of work, and analysis of the Carnot heat engine provides the necessary relationships. The work from a heat engine corresponds to the difference between the heat put into the system at the high temperature, q_{H} and the heat ejected at the low temperature, q_{C}. The efficiency is the work divided by the heat put into the system or:
where w_{cy} is the work done per cycle. The efficiency depends only on q_{C}/q_{H}. Because q_{C} and q_{H} correspond to heat transfer at the temperatures T_{C} and T_{H}, respectively, q_{C}/q_{H} should be some function of these temperatures:
Carnot's theorem states that all reversible engines operating between the same heat reservoirs are equally efficient. Thus, a heat engine operating between T_{1} and T_{3} must have the same efficiency as one consisting of two cycles, one between T_{1} and T_{2}, and the second between T_{2} and T_{3}. This can only be the case if:
which implies:
Since the first function is independent of T_{2}, this temperature must cancel on the right side, meaning f(T_{1},T_{3}) is of the form g(T_{1})/g(T_{3}) (i.e. f(T_{1},T_{3}) = f(T_{1},T_{2})f(T_{2},T_{3}) = g(T_{1})/g(T_{2})· g(T_{2})/g(T_{3}) = g(T_{1})/g(T_{3})), where g is a function of a single temperature. A temperature scale can now be chosen with the property that:
Substituting Equation 4 back into Equation 2 gives a relationship for the efficiency in terms of temperature:
For T_{C} = 0 K the efficiency is 100% and that efficiency becomes greater than 100% below 0 K. Since an efficiency greater than 100% violates the first law of thermodynamics, this implies that 0 K is the minimum possible temperature. In fact the lowest temperature ever obtained in a macroscopic system was 20 nK, which was achieved in 1995 at NIST. Subtracting the right hand side of Equation 5 from the middle portion and rearranging gives:
where the negative sign indicates heat ejected from the system. This relationship suggests the existence of a state function, S, defined by:
where the subscript indicates a reversible process. The change of this state function around any cycle is zero, as is necessary for any state function. This function corresponds to the entropy of the system, which was described previously. Rearranging Equation 6 gives a new definition for temperature in terms of entropy and heat:
For a system, where entropy S(E) is a function of its energy E, the temperature T is given by:
i.e. the reciprocal of the temperature is the rate of increase of entropy with respect to energy.
Definition from statistical mechanics
Statistical mechanics defines temperature based on a system's fundamental degrees of freedom. Eq.(10) is the defining relation of temperature. Eq. (9) can be derived from the principles underlying the fundamental thermodynamic relation.
Generalized temperature from single particle statistics
It is possible to extend the definition of temperature even to systems of few particles, like in a quantum dot. The generalized temperature is obtained by considering time ensembles instead of configuration space ensembles given in statistical mechanics in the case of thermal and particle exchange between a small system of fermions (N even less than 10) with a single/double occupancy system. The finite quantum grand canonical ensemble,^{[67]} obtained under the hypothesis of ergodicity and orthodicity,^{[68]} allows to express the generalized temperature from the ratio of the average time of occupation _{1} and _{2} of the single/double occupancy system:^{[69]}
where E_{F} is the Fermi energy which tends to the ordinary temperature when N goes to infinity.
Negative temperature
{{#invoke:mainmain}} On the empirical temperature scales, which are not referenced to absolute zero, a negative temperature is one below the zeropoint of the scale used. For example, dry ice has a sublimation temperature of Template:Gaps which is equivalent to Template:Gaps. On the absolute Kelvin scale, however, this temperature is 194.6 K. On the absolute scale of thermodynamic temperature no material can exhibit a temperature smaller than or equal to 0 K, both of which are forbidden by the third law of thermodynamics.
In the quantum mechanical description of electron and nuclear spin systems that have a limited number of possible states, and therefore a discrete upper limit of energy they can attain, it is possible to obtain a negative temperature, which is numerically indeed less than absolute zero. However, this is not the macroscopic temperature of the material, but instead the temperature of only very specific degrees of freedom, that are isolated from others and do not exchange energy by virtue of the equipartition theorem.
A negative temperature is experimentally achieved with suitable radio frequency techniques that cause a population inversion of spin states from the ground state. As the energy in the system increases upon population of the upper states, the entropy increases as well, as the system becomes less ordered, but attains a maximum value when the spins are evenly distributed among ground and excited states, after which it begins to decrease, once again achieving a state of higher order as the upper states begin to fill exclusively. At the point of maximum entropy, the temperature function shows the behavior of a singularity, because the slope of the entropy function decreases to zero at first and then turns negative. Since temperature is the inverse of the derivative of the entropy, the temperature formally goes to infinity at this point, and switches to negative infinity as the slope turns negative. At energies higher than this point, the spin degree of freedom therefore exhibits formally a negative thermodynamic temperature. As the energy increases further by continued population of the excited state, the negative temperature approaches zero asymptotically.^{[70]} As the energy of the system increases in the population inversion, a system with a negative temperature is not colder than absolute zero, but rather it has a higher energy than at positive temperature, and may be said to be in fact hotter at negative temperatures. When brought into contact with a system at a positive temperature, energy will be transferred from the negative temperature regime to the positive temperature region.
Examples of temperature
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Temperature  Peak emittance wavelength^{[71]} of blackbody radiation  

Kelvin  Degrees Celsius  
Absolute zero (precisely by definition) 
0 K  −273.15 °C  cannot be defined 
Coldest temperature achieved^{[72]} 
100 pK  Template:Gaps °C  29,000 km 
Coldest Bose–Einstein condensate^{[73]} 
450 pK  Template:Gaps °C  6,400 km 
One millikelvin (precisely by definition) 
0.001 K  −273.149 °C  Template:Gaps m (radio, FM band)^{[74]} 
Cosmic Microwave Background (2013 measurement) 
2.7260 K  −270.424 °C  0.00106301 m (millimeter wavelength microwave) 
Water's triple point (precisely by definition) 
273.16 K  0.01 °C  10,608.3 nm (long wavelength I.R.) 
Water's boiling point^{[A]}  373.1339 K  99.9839 °C  7,766.03 nm (mid wavelength I.R.) 
Incandescent lamp^{[B]}  2500 K  ≈2,200 °C  1,160 nm (near infrared)^{[C]} 
Sun's visible surface^{[D]}^{[75]}  5,778 K  5,505 °C  501.5 nm (greenblue light) 
Lightning bolt's channel^{[E]} 
28 kK  28,000 °C  100 nm (far ultraviolet light) 
Sun's core^{[E]}  16 MK  16 million °C  0.18 nm (Xrays) 
Thermonuclear weapon (peak temperature)^{[E]}^{[76]} 
350 MK  350 million °C  8.3×10^{−3} nm (gamma rays) 
Sandia National Labs' Z machine^{[E]}^{[77]} 
2 GK  2 billion °C  1.4×10^{−3} nm (gamma rays)^{[F]} 
Core of a highmass star on its last day^{[E]}^{[78]} 
3 GK  3 billion °C  1×10^{−3} nm (gamma rays) 
Merging binary neutron star system^{[E]}^{[79]} 
350 GK  350 billion °C  8×10^{−6} nm (gamma rays) 
Relativistic Heavy Ion Collider^{[E]}^{[80]} 
1 TK  1 trillion °C  3×10^{−6} nm (gamma rays) 
CERN's proton vs nucleus collisions^{[E]}^{[81]} 
10 TK  10 trillion °C  3×10^{−7} nm (gamma rays) 
Universe 5.391×10^{−44} s after the Big Bang^{[E]} 
1.417×10^{32} K  1.417×10^{32} °C  1.616×10^{−27} nm (Planck Length)^{[82]} 
See also
Notes and references
 ↑ Middleton, W.E.K. (1966), pp. 89–105.
 ↑ ^{2.0} ^{2.1} ^{2.2} Truesdell, C.A. (1980), Sections 11 B, 11H, pages 306–310, 320332.
 ↑ Quinn, T.J. (1983).
 ↑ Quinn, T.J. (1983), pp. 61–83.
 ↑ Schooley, J.F. (1986), pp. 115–138.
 ↑ Adkins, C.J. (1968/1983), pp. 119–120.
 ↑ Buchdahl, H.A. (1966), pp. 137–138.
 ↑ Tschoegl, N.W. (2000), p.88.
 ↑ Quinn, T.J. (1983), pp. 98–107.
 ↑ Schooley, J.F. (1986), pp. 138–143.
 ↑ Zeppenfeld, M., Englert, B.G.U., Glöckner, R., Prehn, A., Mielenz, M., Sommer, C., van Buuren, L.D., Motsch, M., Rempe, G. (2012).
 ↑ Miller, J. (2013).
 ↑ Thomson, W. (Lord Kelvin) (1848).
 ↑ Thomson, W. (Lord Kelvin) (1851).
 ↑ Partington, J.R. (1949), pp. 175–177.
 ↑ Roberts, J.K., Miller, A.R. (1928/1960), pp. 321–322.
 ↑ Buchdahl, H.A (1986). On the redundancy of the zeroth law of thermodynamics, J. Phys. A, Math. Gen., 19: L561–L564.
 ↑ {{#invoke:Citation/CS1citation CitationClass=journal }} A partly reliable translation is to be found at Kestin, J. (1976). The Second Law of Thermodynamics, Dowden, Hutchinson & Ross, Stroudsburg PA.
 ↑ Maxwell, J.C. (1871). Theory of Heat, Longmans, Green, and Co., London, p. 57.
 ↑ Bailyn, M. (1994). A Survey of Thermodynamics, American Institute of Physics Press, New York, ISBN 0883187973, page 24.
 ↑ Quinn, T.J. (1983). Temperature, Academic Press, London, ISBN0125696809, pp. 160–162.
 ↑ Tisza, L. (1966). Generalized Thermodynamics, M.I.T. Press, Cambridge MA, pp. 47,57.
 ↑ ^{23.0} ^{23.1} ^{23.2} Münster, A. (1970), Classical Thermodynamics, translated by E.S. Halberstadt, Wiley–Interscience, London, ISBN 0471624306, pp. 49, 69.
 ↑ ^{24.0} ^{24.1} Bailyn, M. (1994). A Survey of Thermodynamics, American Institute of Physics Press, New York, ISBN 0883187973, pp. 14–15, 214.
 ↑ ^{25.0} ^{25.1} Callen, H.B. (1960/1985), Thermodynamics and an Introduction to Thermostatistics, (first edition 1960), second edition 1985, John Wiley & Sons, New York, ISBN 0–471–86256–8, pp. 146–148.
 ↑ Kondepudi, D., Prigogine, I. (1998). Modern Thermodynamics. From Heat Engines to Dissipative Structures, John Wiley, Chichester, ISBN 0471973947, pp. 115–116.
 ↑ Tisza, L. (1966). Generalized Thermodynamics, M.I.T. Press, Cambridge MA, p. 58.
 ↑ Milne, E.A. (1929). The effect of collisions on monochromatic radiative equilibrium, Monthly Notices of the Royal Astronomical Society, 88: 493–502.
 ↑ Gyarmati, I. (1970). Nonequilibrium Thermodynamics. Field Theory and Variational Principles, translated by E. Gyarmati and W.F. Heinz, Springer, Berlin, pp. 63–66
 ↑ Glansdorff, P., Prigogine, I., (1971). Thermodynamic Theory of Structure, Stability and Fluctuations, Wiley, London, ISBN 0471302805, pp. 14–16.
 ↑ Bailyn, M. (1994). A Survey of Thermodynamics, American Institute of Physics Press, New York, ISBN 0883187973, pp. 133–135.
 ↑ Callen, H.B. (1960/1985), Thermodynamics and an Introduction to Thermostatistics, (first edition 1960), second edition 1985, John Wiley & Sons, New York, ISBN 0–471–86256–8, pp. 309–310.
 ↑ Bryan, G.H. (1907). Thermodynamics. An Introductory Treatise dealing mainly with First Principles and their Direct Applications, B.G. Teubner, Leipzig, page 3.[1]
 ↑ ^{34.0} ^{34.1} ^{34.2} ^{34.3} Mach, E. (1900). Die Principien der Wärmelehre. Historischkritisch entwickelt, Johann Ambrosius Barth, Leipzig, section 22, pages 5657.
 ↑ ^{35.0} ^{35.1} Serrin, J. (1986). Chapter 1, 'An Outline of Thermodynamical Structure', pages 332, especially page 6, in New Perspectives in Thermodynamics, edited by J. Serrin, Springer, Berlin, ISBN 3540159312.
 ↑ Maxwell, J.C. (1872). Theory of Heat, third edition, Longmans, Green, London, page 32.
 ↑ Tait, P.G. (1884). Heat, Macmillan, London, Chapter VII, pages 3940.
 ↑ Planck, M. (1897/1903). Treatise on Thermodynamics, translated by A. Ogg, Longmans, Green, London, pages 12.
 ↑ Planck, M. (1914), The Theory of Heat Radiation, second edition, translated into English by M. Masius, Blakiston's Son & Co., Philadelphia, reprinted by Kessinger.
 ↑ {{#invoke:citation/CS1citation CitationClass=book }}
 ↑ {{#invoke:citation/CS1citation CitationClass=book }}
 ↑ {{#invoke:citation/CS1citation CitationClass=book }}
 ↑ Template:Cite web
 ↑ Tait, P.G. (1884). Heat, Macmillan, London, Chapter VII, pages 42, 103117.
 ↑ Beattie, J.A., Oppenheim, I. (1979). Principles of Thermodynamics, Elsevier Scientific Publishing Company, Amsterdam, 0–444–41806–7, page 29.
 ↑ Landsberg, P.T. (1961). Thermodynamics with Quantum Statistical Illustrations, Interscience Publishers, New York, page 17.
 ↑ {{#invoke:Citation/CS1citation CitationClass=journal }}
 ↑ Maxwell, J.C. (1872). Theory of Heat, third edition, Longman's, Green & Co, London, page 45.
 ↑ ^{49.0} ^{49.1} Pitteri, M. (1984). On the axiomatic foundations of temperature, Appendix G6 on pages 522544 of Rational Thermodynamics, C. Truesdell, second edition, Springer, New York, ISBN 0387908749.
 ↑ Truesdell, C., Bharatha, S. (1977). The Concepts and Logic of Classical Thermodynamics as a Theory of Heat Engines, Rigorously Constructed upon the Foundation Laid by S. Carnot and F. Reech, Springer, New York, ISBN 0387079718, page 20.
 ↑ ^{51.0} ^{51.1} Serrin, J. (1978). The concepts of thermodynamics, in Contemporary Developments in Continuum Mechanics and Partial Differential Equations. Proceedings of the International Symposium on Continuum Mechanics and Partial Differential Equations, Rio de Janeiro, August 1977, edited by G.M. de La Penha, L.A.J. Medeiros, NorthHolland, Amsterdam, ISBN 0444851666, pages 411451.
 ↑ Maxwell, J.C. (1872). Theory of Heat, third edition, Longmans, Green, London, pages 155158.
 ↑ Tait, P.G. (1884). Heat, Macmillan, London, Chapter VII, Section 95, pages 6869.
 ↑ Buchdahl, H.A. (1966), p. 73.
 ↑ Kondepudi, D. (2008). Introduction to Modern Thermodynamics, Wiley, Chichester, ISBN 9780470015988, Section 32., pages 106108.
 ↑ The kelvin in the SI Brochure
 ↑ Template:Cite web
 ↑ {{#invoke:Citation/CS1citation CitationClass=journal }}
 ↑ {{#invoke:Citation/CS1citation CitationClass=journal }}
 ↑ Balescu, R. (1975). Equilibrium and Nonequilibrium Statistical Mechanics, Wiley, New York, ISBN 0471046000, pages 148154.
 ↑ ^{61.0} ^{61.1} {{#invoke:citation/CS1citation CitationClass=book }}
 ↑ {{#invoke:Citation/CS1citation CitationClass=journal }}
 ↑ Feynman, R.P., Leighton, R.B., Sands, M. (1963). The Feynman Lectures on Physics, Addison–Wesley, Reading MA, volume 1, pages 39–6 to 39–12.
 ↑ http://galileo.phys.virginia.edu/classes/252/kinetic_theory.html
 ↑ Tolman, R.C. (1938). The Principles of Statistical Mechanics, Oxford University Press, London, pp. 93, 655.
 ↑ {{#invoke:citation/CS1citation CitationClass=book }}
 ↑ {{#invoke:Citation/CS1citation CitationClass=journal }} arxiv.org
 ↑ http://tnt.phys.uniroma1.it/twiki/pub/TNTgroup/AngeloVulpiani/Dellago.pdf
 ↑ {{#invoke:Citation/CS1citation CitationClass=journal }} arxiv.org
 ↑ {{#invoke:citation/CS1citation CitationClass=book }}
 ↑ The cited emission wavelengths are for black bodies in equilibrium. CODATA 2006 recommended value of Template:Val used for Wien displacement law constant b.
 ↑ Template:Cite web
 ↑ A temperature of 450 ±80 pK in a Bose–Einstein condensate (BEC) of sodium atoms was achieved in 2003 by researchers at MIT. Citation: Cooling Bose–Einstein Condensates Below 500 Picokelvin, A. E. Leanhardt et al., Science 301, 12 Sept. 2003, p. 1515. It's noteworthy that this record's peak emittance blackbody wavelength of 6,400 kilometers is roughly the radius of Earth.
 ↑ The peak emittance wavelength of Template:Gaps m is a frequency of 103.456 MHz
 ↑ Measurement was made in 2002 and has an uncertainty of ±3 kelvin. A 1989 measurement produced a value of 5,777.0±2.5 K. Citation: Overview of the Sun (Chapter 1 lecture notes on Solar Physics by Division of Theoretical Physics, Dept. of Physical Sciences, University of Helsinki).
 ↑ The 350 MK value is the maximum peak fusion fuel temperature in a thermonuclear weapon of the Teller–Ulam configuration (commonly known as a hydrogen bomb). Peak temperatures in Gadgetstyle fission bomb cores (commonly known as an atomic bomb) are in the range of 50 to 100 MK. Citation: Nuclear Weapons Frequently Asked Questions, 3.2.5 Matter At High Temperatures. Link to relevant Web page. All referenced data was compiled from publicly available sources.
 ↑ Peak temperature for a bulk quantity of matter was achieved by a pulsedpower machine used in fusion physics experiments. The term bulk quantity draws a distinction from collisions in particle accelerators wherein high temperature applies only to the debris from two subatomic particles or nuclei at any given instant. The >2 GK temperature was achieved over a period of about ten nanoseconds during shot Z1137. In fact, the iron and manganese ions in the plasma averaged 3.58±0.41 GK (309±35 keV) for 3 ns (ns 112 through 115). Ion Viscous Heating in a Magnetohydrodynamically Unstable Z Pinch at Over Template:Val Kelvin, M. G. Haines et al., Physical Review Letters 96 (2006) 075003. Link to Sandia's news release.
 ↑ Core temperature of a high–mass (>8–11 solar masses) star after it leaves the main sequence on the Hertzsprung–Russell diagram and begins the alpha process (which lasts one day) of fusing silicon–28 into heavier elements in the following steps: sulfur–32 → argon–36 → calcium–40 → titanium–44 → chromium–48 → iron–52 → nickel–56. Within minutes of finishing the sequence, the star explodes as a Type II supernova. Citation: Stellar Evolution: The Life and Death of Our Luminous Neighbors (by Arthur Holland and Mark Williams of the University of Michigan). Link to Web site. More informative links can be found here [2], and here [3], and a concise treatise on stars by NASA is here [4]. Template:Dead link
 ↑ Based on a computer model that predicted a peak internal temperature of 30 MeV (350 GK) during the merger of a binary neutron star system (which produces a gamma–ray burst). The neutron stars in the model were 1.2 and 1.6 solar masses respectively, were roughly 20 km in diameter, and were orbiting around their barycenter (common center of mass) at about 390 Hz during the last several milliseconds before they completely merged. The 350 GK portion was a small volume located at the pair's developing common core and varied from roughly 1 to 7 km across over a time span of around 5 ms. Imagine two citysized objects of unimaginable density orbiting each other at the same frequency as the G4 musical note (the 28th white key on a piano). It's also noteworthy that at 350 GK, the average neutron has a vibrational speed of 30% the speed of light and a relativistic mass (m) 5% greater than its rest mass (m_{0}). Torus Formation in Neutron Star Mergers and WellLocalized Short GammaRay Bursts, R. Oechslin et al. of Max Planck Institute for Astrophysics., arXiv:astroph/0507099 v2, 22 Feb. 2006. An html summary.
 ↑ Results of research by Stefan Bathe using the PHENIX detector on the Relativistic Heavy Ion Collider at Brookhaven National Laboratory in Upton, New York, U.S.A. Bathe has studied goldgold, deuterongold, and protonproton collisions to test the theory of quantum chromodynamics, the theory of the strong force that holds atomic nuclei together. Link to news release.
 ↑ How do physicists study particles? by CERN.
 ↑ The Planck frequency equals Template:Val (which is the reciprocal of one Planck time). Photons at the Planck frequency have a wavelength of one Planck length. The Planck temperature of Template:Val equates to a calculated b /T = λ_{max} wavelength of Template:Val. However, the actual peak emittance wavelength quantizes to the Planck length of Template:Val.
Bibliography of cited references
 Adkins, C.J. (1968/1983). Equilibrium Thermodynamics, (1st edition 1968), third edition 1983, Cambridge University Press, Cambridge UK, ISBN 0521254450.
 Buchdahl, H.A. (1966). The Concepts of Classical Thermodynamics, Cambridge University Press, Cambridge UK.
 Middleton, W.E.K. (1966). A History of the Thermometer and its Use in Metrology, Johns Hopkins Press, Baltimore MD.
 Miller, J. (2013). Cooling molecules the optoelectric way, Physics Today, 66(1): 12–14.
 Partington, J.R. (1949). An Advanced Treatise on Physical Chemistry, volume 1, Fundamental Principles. The Properties of Gases, Longmans, Green & Co., London, pp. 175–177.
 Quinn, T.J. (1983). Temperature, Academic Press, London, ISBN 0125696809.
 Schooley, J.F. (1986). Thermometry, CRC Press, Boca Raton, ISBN 0849358337.
 Roberts, J.K., Miller, A.R. (1928/1960). Heat and Thermodynamics, (first edition 1928), fifth edition, Blackie & Son Limited, Glasgow.
 Thomson, W. (Lord Kelvin) (1848). On an absolute thermometric scale founded on Carnot's theory of the motive power of heat, and calculated from Regnault's observations, Proc. Cambridge Phil. Soc. (1843/1863) 1, No. 5: 66–71.
 {{#invoke:Citation/CS1citation
CitationClass=journal }}
 Truesdell, C.A. (1980). The Tragicomical History of Thermodynamics, 18221854, Springer, New York, ISBN 0387904034.
 Tschoegl, N.W. (2000). Fundamentals of Equilibrium and SteadyState Thermodynamics, Elsevier, Amsterdam, ISBN 0444504265.
 Zeppenfeld, M., Englert, B.G.U., Glöckner, R., Prehn, A., Mielenz, M., Sommer, C., van Buuren, L.D., Motsch, M., Rempe, G. (2012). Sysiphus cooling of electrically trapped polyatomic molecules, Nature, 491: 570–573.
Further reading
 Chang, Hasok (2004). Inventing Temperature: Measurement and Scientific Progress. Oxford: Oxford University Press. ISBN 9780195171273.
 Zemansky, Mark Waldo (1964). Temperatures Very Low and Very High. Princeton, N.J.: Van Nostrand.
External links
Template:Sister Template:Sister
 An elementary introduction to temperature aimed at a middle school audience
 from Oklahoma State University
 Average yearly temperature by country A tabular list of countries and Thermal Map displaying the average yearly temperature by country
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