Question:

# What if my body temperature is always warmer than others?

## Body temperature is an individual thing. The average body temp for people is 98.6 degrees, but that is only the average. Some people's body temp is 97 degrees and would be feverish at 98.6.

Hyperthermia is elevated body temperature due to failed thermoregulation that occurs when a body produces or absorbs more heat than it dissipates. Extreme temperature elevation then becomes a medical emergency requiring immediate treatment to prevent disability or death. The most common causes include heat stroke and adverse reactions to drugs. The former is an acute temperature elevation caused by exposure to excessive heat, or combination of heat and humidity, that overwhelms the heat-regulating mechanisms. The latter is a relatively rare side effect of many drugs, particularly those that affect the central nervous system. Malignant hyperthermia is a rare complication of some types of general anesthesia. Hyperthermia can also be deliberately induced using drugs or medical devices and may be used in the treatment of some kinds of cancer and other conditions, most commonly in conjunction with radiotherapy. Hyperthermia differs from fever in that the body's temperature set point remains unchanged. The opposite is hypothermia, which occurs when the temperature drops below that required to maintain normal metabolism. Hyperthermia is defined as a temperature greater than 37.5–38.3 °C (100–101 °F), depending on the reference used, that occurs without a change in the body's temperature set point. The normal human body temperature in health can be as high as in the late afternoon. Hyperthermia requires an elevation from the temperature that would otherwise be expected. Such elevations range from mild to extreme; body temperatures above can be life threatening. Hot, dry, skin is typical as blood vessels dilate in an attempt to increase heat loss. An inability to cool the body through perspiration may cause the skin to feel dry. Other signs and symptoms vary. Accompanying dehydration can produce nausea, vomiting, headaches, and low blood pressure and the latter can lead to fainting or dizziness, especially if the standing position is assumed quickly. In severe heat stroke, there may be confused, hostile, or seemingly intoxicated behavior. Heart rate and respiration rate will increase (tachycardia and tachypnea) as blood pressure drops and the heart attempts to maintain adequate circulation. The decrease in blood pressure can then cause blood vessels to contract reflexly, resulting in a pale or bluish skin color in advanced cases. Young children, in particular, may have seizures. Eventually, organ failure, unconsciousness and death will result. Heat stroke occurs when thermoregulation is overwhelmed by a combination of excessive metabolic production of heat (exertion), excessive environmental heat, and insufficient or impaired heat loss, resulting in an abnormally high body temperature. In severe cases, temperatures can exceed . Heat stroke may be non-exertional (classic) or exertional.
Significant physical exertion in hot conditions can generate heat beyond the ability to cool, because, in addition to the heat, humidity of the environment may reduce the efficiency of the body's normal cooling mechanisms. Heat loss mechanisms are limited to vasodilation of skin vessels and increased rate of sweating. Vasodilation dissipates heat by convection and sweating by evaporation. However, thermoregulation can be assisted with shade or fans. Other factors, such as insufficient water intake, consuming alcohol, or lack of air conditioning, can worsen the problem.
The principles of physics involved include: Non-exertional heat stroke mostly affects the young and elderly. In the elderly in particular, it can be precipitated by medications such as anticholinergic drugs, antihistamines, and diuretics that reduce vasodilation, sweating, and other heat-loss mechanisms. In this situation, the body's tolerance for high environmental temperature may be insufficient, even at rest.

The Köppen climate classification is one of the most widely used climate classification systems. It was first published by Russian German climatologist Wladimir Köppen in 1884, with several later modifications by Köppen himself, notably in 1918 and 1936. Later, German climatologist Rudolf Geiger collaborated with Köppen on changes to the classification system, which is thus sometimes referred to as the Köppen–Geiger climate classification system. The system is based on the concept that native vegetation is the best expression of climate. Thus, climate zone boundaries have been selected with vegetation distribution in mind. It combines average annual and monthly temperatures and precipitation, and the seasonality of precipitation.:200–1 The Köppen climate classification scheme divides climates into five main groups, each having several types and subtypes. Each particular climate type is represented by a 2 to 4 letter symbol. Tropical climates are characterized by constant high temperature (at sea level and low elevations) — all twelve months of the year have average temperatures of or higher. They are subdivided as follows: These climates are characterized by the fact that actual precipitation is less than a threshold value set equal to the potential evapotranspiration.:212 The threshold value (in millimeters) is determined as follows: These climates have an average temperature above in their warmest months (April to September in northern hemisphere), and a coldest month average between −3 and 18 °C (27 and 64 °F). Some climatologists, particularly in the United States, however, prefer to observe rather than in the coldest month as the boundary between this group and the colder Group D (Humid Continental).This is also done to prevent certain mild headland locations on the upper East Coast of the USA and Japan from fitting into the C group. When the boundary between C (Mild Temperate/mesothermal climates) and D (Cold winter/microthermal climates) is increased to 32 F (not the 27 F suggested by Köppen), this creates a smaller C zone located further southward. In the USA, areas from the NYC metropolitan area (NYC/New Jersey/southern Connecticut) southward, as well as the lower Ohio Valley, lower Midwest, and southern Plains are located in the mild C group...while locations to the north of these regions (Northern Plains, Great Lakes, Midwest, upper Ohio Valley and upper East Coast (Boston northward), are located in the cooler D group. Using 32 F also pushes parts of the northeast and northcentral Asia (northern Japan, northern China, and northern Korea) into the colder D/microthermal group (aka known as humid continental).
These climates have an average temperature above in their warmest months, and a coldest month average below −3 °C (or 0 °C in some versions, as noted previously). These usually occur in the interiors of continents and on their upper east coasts, normally north of 40° North latitude. In the Southern Hemisphere, Group D climates are extremely rare due to the smaller land masses in the middle latitudes and the almost complete absence of land at 40°–60° South latitude, existing only in some highland locations. Group D climates are subdivided as follows: Lettering Scheme These climates are characterized by average temperatures below in all twelve months of the year: Some climatologists have argued that Köppen's system could be improved upon. One of the most frequently-raised objections concerns the temperate Group C category, regarded by many as overbroad. Using the 0°C isotherm, New Orleans, LA and London would both fall into this climate scheme, despite dramatic differences between these 2 locations. In Applied Climatology (first edition published in 1966), John F. Griffiths proposed a new subtropical zone, encompassing those areas with a coldest month of between 6 and 18 °C (43 and 64 °F), effectively subdividing Group C into two nearly equal parts (his scheme assigns the letter B to the new zone, and identifies dry climates with an additional letter immediately following the temperature-based letter). Another point of contention involves the dry B climates; the argument here is that their separation by Köppen into only two thermal subsets is inadequate. Those who hold this view (including Griffiths) have suggested that the dry climates be placed on the same temperature continuum as other climates, with the thermal letter being followed by an additional capital letter — S for steppe or W (or D) for desert — as applicable (Griffiths also advances an alternate formula for use as an aridity threshold: R = 160 + 9T, with R equalling the threshold, in millimeters of mean annual precipitation, and T denoting the mean annual temperature in degrees Celsius). A third idea is to create a maritime polar or EM zone within Group E to separate relatively mild marine locations (such as the Falkland Islands, and the outer Aleutian Islands) from the colder, continental tundra climates. Specific proposals vary; some advocate setting a coldest-month parameter, such as , while others support assigning the new designation to areas with an average annual temperature of above 0 °C. The accuracy of the 10 °C warmest-month line as the start of the polar climates has also been questioned; Otto Nordenskiöld, for example, devised an alternate formula: W = 9 − 0.1 C, with W representing the average temperature of the warmest month and C that of the coldest month, both in degrees Celsius (for instance, if the coldest month averaged −20 °C, a warmest-month average of 11 °C or higher would be necessary to prevent the climate from being polar). This boundary does appear to more closely follow the tree line, or the latitude poleward of which trees cannot grow, than the 10 °C warmest-month isotherm; the former tends to run poleward of the latter near the western margins of the continents, but at a lower latitude in the landmass interiors, the two lines crossing at or near the east coasts of both Asia and North America. The Trewartha climate classification scheme (1966 and 1980 update) is a modified version of the Köppen system, and was an answer to some of the deficiencies of the 1899 Köppen system. The newer Trewartha theme attempts to redefine the middle latitudes in such a way as to be closer to vegetational zoning and genetic climate systems. This change was seen as most effective in Asia and North America, where many areas fell into a single zone (the C climate group). Under the standard Köppen system in the USA for example, western Washington and Oregon are classed into the same climate as southern California, even though the two regions have strikingly different weather and vegetation. The Köppen system also classes Midwest into the same climate as the Gulf Coast. Trewartha's modifications sought to reclass the middle latitudes into zones; 1) Subtropical - 8 or more months have a mean temperature of 50 F/10 C or higher. 2) Temperate - 4 to 7 months have a mean temperature of 10 C or higher. 3) Boreal (or subarctic) - 1 to 3 months have a mean temperature of 10 C or higher. This change from the older Köppen system was thought to reflect a more true or "real world" reflection of the global climate. Based on recent data sets from the Climatic Research Unit (CRU) of the University of East Anglia and the Global Precipitation Climatology Centre (GPCC) at the German Weather Service, a new digital Köppen–Geiger world map on climate classification for the second half of the 20th century has been compiled. All maps use the ≥0 °C definition for temperate climates and the 18 °C annual mean temperature threshold to distinguish between hot and cold dry climates. Köppen map of Africa Köppen map of the Americas Köppen map of Asia Köppen map of Australia/Oceania Köppen map of Brazil Köppen map of Europe Köppen map of North America Köppen map of South Asia Köppen map of Russia Köppen map of South America Köppen map of the Middle East

Atmospheric temperature range is the numerical difference between the minimum and maximum values of temperature observed in a given location. A temperature range may refer to a period of time (e.g., in a given day, month, year, century) or to an average (average of all temperature ranges in a period of time). The variation in temperature that occurs from the highs of the day to the cool of nights is called diurnal temperature variation. The size of ground-level atmospheric temperature ranges depends on several factors, such as: A location which combines an average temperature of 19 degrees Celsius, 60% average humidity and a temperature range of about 10 degrees Celsius around the average temperature (yearly temperature variation) is considered ideal in terms of comfort for the human species. Most of the places with these characteristics are located in the transition between temperate and tropical climates, approximately around the tropics, particularly in the Southern hemisphere (the tropic of Capricorn). The figure at left shows an example of monthly temperatures recorded at one of such locations, the city of Campinas, state of São Paulo, Brazil, which lies approximately 60 km north of the Capricorn line (latitude of 22 degrees). Average yearly temperature is 22.4 degrees Celsius, ranging from an average minimum of 12.2 degrees to a maximum of 29.9 degrees. The average temperature range is 11.4 degrees [1]. Variability along the year is small (standard deviation of 2.31 for the maximum monthly average and 4.11 for the minimum). One can easily see in the graph another typical phenomenon of temperature ranges, which is its increase during winter (lower average air temperature). In Campinas, for example, the daily temperature range in July (the coolest month of the year) may vary between typically 10 and 24 degrees Celsius (range of 14), while in January, it may range between 20 and 30 degrees Celsius (range of 10). The effect of latitude, tropical climate, constant gentle wind and sea-side locations clearly show smaller average temperature ranges, smaller variations of temperature, and a higher average temperature (second graph, taken for the same period as Campinas, at Aracaju, capital of the state of Sergipe, also in Brazil, at a latitude of 10 degrees, nearer to the Equator). Average maximum yearly temperature is 28.7 degrees Celsius and average minimum is 21.9. The average temperature range is 5.7 degrees only. Temperature variation along the year in Aracaju is very damped (standard deviation of 1.93 for the maximum temperature and 2.72 for the minimum temperature). The minimum temperature at night does not occur on the ground but few tens of centimeters above the ground. The lowest temperature layer is called Ramdas layer after L. A. Ramdas, who first reported this phenomenon in 1932 based on observations at different screen heights at six meteorological centers across India. The phenomenon is attributed to the interaction of thermal radiation effects on atmospheric aerosols and convection transfer close to the ground.

Molecules, such as oxygen (O2), 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 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 a higher energy input to increase its temperature by a certain amount, i.e. it will have a higher heat capacity than a monatomic gas. The process of cooling involves removing thermal energy from a system. When no more energy can be removed, the system is at absolute zero, which 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 motion of the particles comprising matter would cease and they would be at complete rest in this classical sense. Microscopically in the description of quantum mechanics, however, matter still has zero-point energy even at absolute zero, because of the uncertainty principle. Temperature is a measure of a quality of a state of a material 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 well-defined hotness or temperature. Hotness may be represented abstractly as a one-dimensional manifold. Every valid temperature scale has its own one-to-one map into the hotness manifold. 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. Experimental physicists, for example Galileo and Newton, found that there are indefinitely many empirical temperature scales. Nevertheless, the zeroth law of thermodynamics says that they all measure the same quality. 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. 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. 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 one-dimensional manifold. This is a fundamental character of temperature and thermometers for bodies in their own thermodynamic equilibrium. Except for a system undergoing a first-order 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. 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, non-absolute, hotness and temperature, for a suitable range of processes. This is a matter for study in non-equilibrium thermodynamics. 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 non-equilibrium thermodynamics. 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. 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. 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. When a sample is heated, meaning it receives thermal energy from an external source, some of the introduced heat is converted into kinetic energy, the rest to other forms of internal energy, specific to the material. The amount converted into kinetic energy causes the temperature of the material to rise. The introduced heat ($\Delta Q$) divided by the observed temperature change is the heat capacity (C) of the material. 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 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 non-scientific 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 = , or absolute zero. Many engineering fields in the U.S., notably high-tech 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 thermodynamic-related disciplines such as combustion. 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 corresponds very closely to the freezing point of water and is its boiling point at sea level. Because liquid droplets commonly exist in clouds at sub-zero temperatures, is better defined as the melting point of ice. In this scale a temperature difference of 1 degree Celsius is the same as a increment, but the scale is offset by the temperature at which ice melts (273.15 K). By international agreement 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 and . 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. Quantum-mechanically, however, zero-point motion remains and has an associated energy, the zero-point energy. Matter is in its ground state, and contains no thermal energy. The triple point of water is defined as and . 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 ( = and = ). 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. The following table shows the temperature conversion formulas for conversions to and from the Celsius scale. The field of plasma physics deals with phenomena of electromagnetic nature that involve very high temperatures. It is customary to express temperature in electronvolts (eV) or kiloelectronvolts (keV), where 1 eV = . In the study of QCD matter one routinely encounters temperatures of the order of a few hundred MeV, equivalent to about . 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. 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 quantum-mechanical 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 monoatomic 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. The kinetic theory of gases uses the model of the ideal gas to relate temperature to the average translational kinetic energy of the molecules in a container of gas in thermodynamic equilibrium. Classical mechanics defines the translational kinetic energy of a gas molecule as follows: where m is the particle mass and v its speed, the magnitude of its velocity. The distribution of the speeds (which determine the translational kinetic energies) of the particles in a classical ideal gas is called the Maxwell-Boltzmann distribution. The temperature of a classical ideal gas is related to its average kinetic energy per degree of freedom via the equation: where the Boltzmann constant $k = R/n$ (n = Avogadro number, R = ideal gas constant). This relation is valid in the ideal gas regime, i.e. when the particle density is much less than $1/\Lambda^{3}$, where $\Lambda$ is the thermal de Broglie wavelength. A monoatomic gas has only the three translational degrees of freedom. The zeroth law of thermodynamics implies that any two given systems in thermal equilibrium have the same temperature. In statistical thermodynamics, it can be deduced from the second law of thermodynamics that they also have the same average kinetic energy per particle. In a mixture of particles of various masses, lighter particles move faster than do heavier particles, but have the same average kinetic energy. A neon atom moves slowly relative to a hydrogen molecule of the same kinetic energy. A pollen particle suspended in water moves in a slow Brownian motion among fast-moving water molecules. It has long been recognized that if two bodies of different temperatures are brought into thermal connection, conductive or radiative, they exchange heat accompanied by changes of other state variables. Left isolated from other bodies, the two connected bodies eventually reach a state of thermal equilibrium in which no further changes occur. This basic knowledge is relevant to thermodynamics. Some approaches to thermodynamics take this basic knowledge as axiomatic, other approaches select only one narrow aspect of this basic knowledge as axiomatic, and use other axioms to justify and express deductively the remaining aspects of it. The one aspect chosen by the latter approaches is often stated in textbooks as the zeroth law of thermodynamics, but other statements of this basic knowledge are made by various writers. The usual textbook 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 is taken to justify a statement that all three systems have the same temperature, but, by itself, it does not justify the idea of temperature as a numerical scale for a concept of hotness which exists on a one-dimensional manifold with a sense of greater hotness. Sometimes the zeroth law is stated to provide the latter justification. For suitable systems, an empirical temperature scale may be defined by the variation of one of the other state variables, such as pressure, when all other coordinates are fixed. The second law of thermodynamics is used to define an absolute thermodynamic temperature scale for systems in thermal equilibrium. A temperature scale is based on the properties of some reference system to which other thermometers may be calibrated. One such reference system is a fixed quantity of gas. The ideal gas law indicates that the product of the pressure (p) and volume (V) of a gas is directly proportional to the thermodynamic temperature: where T is temperature, n is the number of moles of gas and R = is the gas constant. Reformulating the pressure-volume term as the sum of classical mechanical particle energies in terms of particle mass, m, and root-mean-square particle speed v, the ideal gas law directly provides the relationship between kinetic energy and temperature: Thus, one can define a scale for temperature based on the corresponding pressure and volume of the gas: the temperature in kelvins is the pressure in pascals of one mole of gas in a container of one cubic metre, divided by the gas constant. In practice, such a gas thermometer is not very convenient, but other thermometers can be calibrated to this scale. The pressure, volume, and the number of moles of a substance are all inherently greater than or equal to zero, suggesting that temperature must also be greater than or equal to zero. As a practical matter it is not possible to use a gas thermometer to measure absolute zero temperature since the gasses 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. 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, qH and the heat ejected at the low temperature, qC. The efficiency is the work divided by the heat put into the system or: where wcy is the work done per cycle. The efficiency depends only on qC/qH. Because qC and qH correspond to heat transfer at the temperatures TC and TH, respectively, qC/qH 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 T1 and T3 must have the same efficiency as one consisting of two cycles, one between T1 and T2, and the second between T2 and T3. This can only be the case if: which implies: Since the first function is independent of T2, this temperature must cancel on the right side, meaning f(T1,T3) is of the form g(T1)/g(T3) (i.e. f(T1,T3) = f(T1,T2)f(T2,T3) = g(T1)/g(T2g(T2)/g(T3) = g(T1)/g(T3)), 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: Notice that for TC = 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. Statistical mechanics defines temperature based on a system's fundamental degrees of freedom. Eq.(8) is the defining relation of temperature. Eq. (7) can be derived from the principles underlying the fundamental thermodynamic relation. 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, obtained under the hypothesis of ergodicity and orthodicity, allows to express the generalized temperature from the ratio of the average time of occupation $\tau$1 and $\tau$2 of the single/double occupancy system: where EF is the Fermi energy which tends to the ordinary temperature when N goes to infinity. On the empirical temperature scales, which are not referenced to absolute zero, a negative temperature is one below the zero-point of the scale used. For example, dry ice has a sublimation temperature of which is equivalent to . 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. 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.

POAH is an acronym for preoptic anterior hypothalamus, the part of the brain that senses core body temperature and regulates it to about 36.8 °C (98.6 °F). -- Method for Heating the Preoptic Anterior Hypothalamus