1.
Energy
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In physics, energy is the property that must be transferred to an object in order to perform work on – or to heat – the object, and can be converted in form, but not created or destroyed. The SI unit of energy is the joule, which is the transferred to an object by the mechanical work of moving it a distance of 1 metre against a force of 1 newton. Mass and energy are closely related, for example, with a sensitive enough scale, one could measure an increase in mass after heating an object. Living organisms require available energy to stay alive, such as the humans get from food. Civilisation gets the energy it needs from energy resources such as fuels, nuclear fuel. The processes of Earths climate and ecosystem are driven by the radiant energy Earth receives from the sun, the total energy of a system can be subdivided and classified in various ways. It may also be convenient to distinguish gravitational energy, thermal energy, several types of energy, electric energy. Many of these overlap, for instance, thermal energy usually consists partly of kinetic. Some types of energy are a mix of both potential and kinetic energy. An example is energy which is the sum of kinetic. Whenever physical scientists discover that a phenomenon appears to violate the law of energy conservation. Heat and work are special cases in that they are not properties of systems, in general we cannot measure how much heat or work are present in an object, but rather only how much energy is transferred among objects in certain ways during the occurrence of a given process. Heat and work are measured as positive or negative depending on which side of the transfer we view them from, the distinctions between different kinds of energy is not always clear-cut. In contrast to the definition, energeia was a qualitative philosophical concept, broad enough to include ideas such as happiness. The modern analog of this property, kinetic energy, differs from vis viva only by a factor of two, in 1807, Thomas Young was possibly the first to use the term energy instead of vis viva, in its modern sense. Gustave-Gaspard Coriolis described kinetic energy in 1829 in its modern sense, the law of conservation of energy was also first postulated in the early 19th century, and applies to any isolated system. It was argued for years whether heat was a physical substance, dubbed the caloric, or merely a physical quantity. In 1845 James Prescott Joule discovered the link between mechanical work and the generation of heat and these developments led to the theory of conservation of energy, formalized largely by William Thomson as the field of thermodynamics

2.
Enthalpy
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Enthalpy /ˈɛnθəlpi/ is a measurement of energy in a thermodynamic system. It is the thermodynamic quantity equivalent to the heat content of a system. It is equal to the energy of the system plus the product of pressure. Enthalpy is defined as a function that depends only on the prevailing equilibrium state identified by the systems internal energy, pressure. The unit of measurement for enthalpy in the International System of Units is the joule, but other historical, conventional units are still in use, such as the British thermal unit and the calorie. At constant pressure, the enthalpy change equals the energy transferred from the environment through heating or work other than expansion work, the total enthalpy, H, of a system cannot be measured directly. The same situation exists in classical mechanics, only a change or difference in energy carries physical meaning. Enthalpy itself is a potential, so in order to measure the enthalpy of a system, we must refer to a defined reference point, therefore what we measure is the change in enthalpy. The ΔH is a change in endothermic reactions, and negative in heat-releasing exothermic processes. For processes under constant pressure, ΔH is equal to the change in the energy of the system. This means that the change in enthalpy under such conditions is the heat absorbed by the material through a reaction or by external heat transfer. Enthalpies for chemical substances at constant pressure assume standard state, most commonly 1 bar pressure, standard state does not, strictly speaking, specify a temperature, but expressions for enthalpy generally reference the standard heat of formation at 25 °C. Enthalpy of ideal gases and incompressible solids and liquids does not depend on pressure, unlike entropy, real materials at common temperatures and pressures usually closely approximate this behavior, which greatly simplifies enthalpy calculation and use in practical designs and analyses. The word enthalpy stems from the Ancient Greek verb enthalpein, which means to warm in and it combines the Classical Greek prefix ἐν- en-, meaning to put into, and the verb θάλπειν thalpein, meaning to heat. The word enthalpy is often attributed to Benoît Paul Émile Clapeyron. This misconception was popularized by the 1927 publication of The Mollier Steam Tables, however, neither the concept, the word, nor the symbol for enthalpy existed until well after Clapeyrons death. The earliest writings to contain the concept of enthalpy did not appear until 1875, however, Gibbs did not use the word enthalpy in his writings. The actual word first appears in the literature in a 1909 publication by J. P. Dalton

3.
Entropy
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In statistical thermodynamics, entropy is a measure of the number of microscopic configurations Ω that a thermodynamic system can have when in a state as specified by some macroscopic variables. Formally, S = k B ln Ω, for example, gas in a container with known volume, pressure, and temperature could have an enormous number of possible configurations of the collection of individual gas molecules. Each instantaneous configuration of the gas may be regarded as random, Entropy may be understood as a measure of disorder within a macroscopic system. The second law of thermodynamics states that an isolated systems entropy never decreases, such systems spontaneously evolve towards thermodynamic equilibrium, the state with maximum entropy. Non-isolated systems may lose entropy, provided their environments entropy increases by at least that amount, since entropy is a function of the state of the system, a change in entropy of a system is determined by its initial and final states. This applies whether the process is reversible or irreversible, however, irreversible processes increase the combined entropy of the system and its environment. The above definition is called the macroscopic definition of entropy because it can be used without regard to any microscopic description of the contents of a system. The concept of entropy has found to be generally useful and has several other formulations. Entropy was discovered when it was noticed to be a quantity that behaves as a function of state and it has the dimension of energy divided by temperature, which has a unit of joules per kelvin in the International System of Units. But the entropy of a substance is usually given as an intensive property—either entropy per unit mass or entropy per unit amount of substance. In statistical mechanics this reflects that the state of a system is generally non-degenerate. Understanding the role of entropy in various processes requires an understanding of how. It is often said that entropy is an expression of the disorder, or randomness of a system, the second law is now often seen as an expression of the fundamental postulate of statistical mechanics through the modern definition of entropy. In other words, in any natural process there exists an inherent tendency towards the dissipation of useful energy and he made the analogy with that of how water falls in a water wheel. This was an insight into the second law of thermodynamics. g. Clausius described entropy as the transformation-content, i. e. dissipative energy use and this was in contrast to earlier views, based on the theories of Isaac Newton, that heat was an indestructible particle that had mass. Later, scientists such as Ludwig Boltzmann, Josiah Willard Gibbs, henceforth, the essential problem in statistical thermodynamics, i. e. according to Erwin Schrödinger, has been to determine the distribution of a given amount of energy E over N identical systems. Carathéodory linked entropy with a definition of irreversibility, in terms of trajectories

4.
Entropy (order and disorder)
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In thermodynamics, entropy is commonly associated with the amount of order, disorder, or chaos in a thermodynamic system. In recent years, in chemistry textbooks there has been a shift away from using the order and disorder to that of the concept of energy dispersion to describe entropy. Locally, the entropy can be lowered by external action and this applies to machines, such as a refrigerator, where the entropy in the cold chamber is being reduced, and to living organisms. This local decrease in entropy is, however, only possible at the expense of an increase in the surroundings. This was the first-ever statistical law in physics, in 1864, Ludwig Boltzmann, a young student in Vienna, came across Maxwell’s paper and was so inspired by it that he spent much of his long and distinguished life developing the subject further. Similarly, in 1882 Hermann von Helmholtz used the word Unordnung to describe entropy, a measure of disorder, the higher the entropy the greater the disorder. In thermodynamics, a representing the state of disorder of a system at the atomic, ionic, or molecular level. A measure of disorder in the universe or of the availability of the energy in a system to do work, Entropy and disorder also have associations with equilibrium. Technically, entropy, from this perspective, is defined as a property which serves as a measure of how close a system is to equilibrium — that is. Likewise, the value of the entropy of a distribution of atoms and molecules in a system is a measure of the disorder in the arrangements of its particles. ”In particular, many biologists have taken to speaking in terms of the entropy of an organism, or about its antonym negentropy. As an example, consider a box that is divided two sections. What is the probability that a number, or all of the particles. If you only have one particle, then that system of one particle can subsist in two states, one side of the box versus the other. If you have more than one particle, or define states as being further locational subdivisions of the box, the laws of thermodynamics seem to dictate the opposite, that nature should inexorably degenerate toward a state of greater disorder, greater entropy. Yet all around us we see magnificent structures—galaxies, cells, ecosystems, human beings—that have all managed to assemble themselves. ”This local increase in order is, however, only possible at the expense of an entropy increase in the surroundings. The conditioner of this statement suffices that living systems are systems in which both heat, mass, and or work may transfer into or out of the system. Unlike temperature, the entropy of a living system would drastically change if the organism were thermodynamically isolated. If an organism was in this type of situation, its entropy would increase markedly as the once-living components of the organism decayed to an unrecognizable mass

5.
Fugacity
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In chemical thermodynamics, the fugacity of a real gas is an effective partial pressure which replaces the mechanical partial pressure in an accurate computation of the chemical equilibrium constant. It is equal to the pressure of a gas which has the same chemical potential as the real gas. For example, nitrogen gas at 0 °C and a pressure of P =100 atm has a fugacity of f =97.03 atm.03 atm. Fugacities are determined experimentally or estimated from various models such as a Van der Waals gas that are closer to reality than an ideal gas, the ideal gas pressure and fugacity are related through the dimensionless fugacity coefficient φ. φ = f P For nitrogen at 100 atm, the fugacity coefficient is 97.03 atm /100 atm =0.9703, for an ideal gas, fugacity and pressure are equal so φ is 1. The contribution of nonideality to the potential of a real gas is equal to RT ln φ. Again for nitrogen at 100 atm, the potential is μ = μid + RT ln 0.9703. The fugacity is closely related to the thermodynamic activity, for a gas, the activity is simply the fugacity divided by a reference pressure to give a dimensionless quantity. Accurate calculations of chemical equilibrium for real gases should use the fugacity rather than the pressure, the thermodynamic condition for chemical equilibrium is that the total chemical potential of reactants is equal to that of products. For a condensed phase in equilibrium with its phase, the chemical potential is equal to that of the vapor. This fugacity is equal to the vapor pressure when the vapor pressure is not too high. The word fugacity is derived from the Latin for fleetness, which is interpreted as the tendency to flee or escape. The concept of fugacity was introduced by American chemist Gilbert N. Lewis in 1901, the fugacity of a real gas is formally defined by an equation analogous to the relation between the chemical potential and the pressure of an ideal gas. In general, the chemical potential μ is defined as the partial molar Gibbs free energy. However, for any pure substance it is equal to the molar Gibbs free energy, at constant temperature, this expression can be integrated as a function of P. We must also set a reference state, for a real gas, the integral ∫ P ∘ P V ¯ d P cannot be calculated because there is no simple expression for a real gass molar volume. Additionally, chemical potential is not mathematically well behaved and it approaches negative infinity as pressure approaches zero and this creates problems in doing real calculations. It is desirable that the expression for a real gass chemical potential to be similar to the one for an ideal gas

6.
Gibbs free energy
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Just as in mechanics, where the decrease in potential energy is defined as maximum useful work that can be performed, similarly different potentials have different meanings. The Gibbs energy is also the potential that is minimized when a system reaches chemical equilibrium at constant pressure and temperature. Its derivative with respect to the coordinate of the system vanishes at the equilibrium point. As such, a reduction in G is a condition for the spontaneity of processes at constant pressure and temperature. The Gibbs free energy, originally called available energy, was developed in the 1870s by the American scientist Josiah Willard Gibbs. The initial state of the body, according to Gibbs, is supposed to be such that the body can be made to pass from it to states of dissipated energy by reversible processes. In his 1876 magnum opus On the Equilibrium of Heterogeneous Substances, according to the second law of thermodynamics, for systems reacting at STP, there is a general natural tendency to achieve a minimum of the Gibbs free energy. A quantitative measure of the favorability of a reaction at constant temperature and pressure is the change ΔG in Gibbs free energy that is caused by the reaction. As a necessary condition for the reaction to occur at constant temperature and pressure, ΔG must be smaller than the non-PV work, ΔG equals the maximum amount of non-PV work that can be performed as a result of the chemical reaction for the case of reversible process. The equation can be seen from the perspective of the system taken together with its surroundings. First assume that the reaction at constant temperature and pressure is the only one that is occurring. Then the entropy released or absorbed by the system equals the entropy that the environment must absorb or release, the reaction will only be allowed if the total entropy change of the universe is zero or positive. This is reflected in a negative ΔG, and the reaction is called exergonic, if we couple reactions, then an otherwise endergonic chemical reaction can be made to happen. In traditional use, the term free was included in Gibbs free energy to mean available in the form of useful work, the characterization becomes more precise if we add the qualification that it is the energy available for non-volume work. However, a number of books and journal articles do not include the attachment free. This is the result of a 1988 IUPAC meeting to set unified terminologies for the scientific community. This standard, however, has not yet been universally adopted. Further, Gibbs stated, In this description, as used by Gibbs, ε refers to the energy of the body, η refers to the entropy of the body

7.
Helmholtz free energy
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In thermodynamics, the Helmholtz free energy is a thermodynamic potential that measures the “useful” work obtainable from a closed thermodynamic system at a constant temperature and volume. The negative of the difference in the Helmholtz energy is equal to the amount of work that the system can perform in a thermodynamic process in which volume is held constant. If the volume is not held constant, part of work will be performed as boundary work. The Helmholtz energy is used for systems held at constant volume. Since in this case no work is performed on the environment, for a system at constant temperature and volume, the Helmholtz energy is minimized at equilibrium. The Helmholtz free energy was developed by Hermann von Helmholtz, a German physician and physicist, the IUPAC recommends the letter A as well as the use of name Helmholtz energy. In physics, the letter F can also be used to denote the Helmholtz energy, as Helmholtz energy is referred to as the Helmholtz function, Helmholtz free energy. For example, in research, Helmholtz free energy is often used since explosive reactions by their nature induce pressure changes. It is also used to define fundamental equations of state of pure substances. The Helmholtz energy is the Legendre transform of the energy, U. If the system is kept at fixed volume and is in contact with a bath at some constant temperature. Conservation of energy implies, Δ U bath + Δ U + W =0 The volume of the system is kept constant and this means that the volume of the heat bath does not change either and we can conclude that the heat bath does not perform any work. This implies that the amount of heat that flows into the bath is given by. This result seems to contradict the equation dA = -S dT - P dV, as keeping T and V constant seems to imply dA =0, to allow for spontaneous processes at constant T and V, one needs to enlarge the thermodynamical state space of the system. In case of a reaction, one must allow for changes in the numbers Nj of particles of each type j. This equation is then valid for both reversible and non-reversible uPT changes. In case of a change at constant T and V without electrical work. A system kept at constant volume, temperature, and particle number is described by the canonical ensemble, the fact that the system does not have a unique energy means that the various thermodynamical quantities must be defined as expectation values

8.
Internal energy
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It keeps account of the gains and losses of energy of the system that are due to changes in its internal state. The internal energy of a system can be changed by transfers of matter or heat or by doing work, when matter transfer is prevented by impermeable containing walls, the system is said to be closed. Then the first law of thermodynamics states that the increase in energy is equal to the total heat added plus the work done on the system by its surroundings. If the containing walls pass neither matter nor energy, the system is said to be isolated, the first law of thermodynamics may be regarded as establishing the existence of the internal energy. The internal energy is one of the two cardinal state functions of the variables of a thermodynamic system. The internal energy of a state of a system cannot be directly measured. Such a chain, or path, can be described by certain extensive state variables of the system, namely, its entropy, S, its volume, V. The internal energy, U, is a function of those, sometimes, to that list are appended other extensive state variables, for example electric dipole moment. Customarily, thermodynamic descriptions include only items relevant to the processes under study, Thermodynamics is chiefly concerned only with changes in the internal energy, not with its absolute value. The internal energy is a function of a system, because its value depends only on the current state of the system. It is the one and only cardinal thermodynamic potential, through it, by use of Legendre transforms, are mathematically constructed the other thermodynamic potentials. These are functions of variable lists in which some extensive variables are replaced by their conjugate intensive variables, Legendre transformation is necessary because mere substitutive replacement of extensive variables by intensive variables does not lead to thermodynamic potentials. Mere substitution leads to a less informative formula, an equation of state, though it is a macroscopic quantity, internal energy can be explained in microscopic terms by two theoretical virtual components. One is the kinetic energy due to the microscopic motion of the systems particles. The other is the energy associated with the microscopic forces, including the chemical bonds. If thermonuclear reactions are specified as a topic of concern, then the static rest mass energy of the constituents of matter is also counted. There is no simple relation between these quantities of microscopic energy and the quantities of energy gained or lost by the system in work, heat. The SI unit of energy is the joule, sometimes it is convenient to use a corresponding density called specific internal energy which is internal energy per unit of mass of the system in question

9.
Pressure
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Pressure is the force applied perpendicular to the surface of an object per unit area over which that force is distributed. Gauge pressure is the relative to the ambient pressure. Various units are used to express pressure, Pressure may also be expressed in terms of standard atmospheric pressure, the atmosphere is equal to this pressure and the torr is defined as 1⁄760 of this. Manometric units such as the centimetre of water, millimetre of mercury, Pressure is the amount of force acting per unit area. The symbol for it is p or P, the IUPAC recommendation for pressure is a lower-case p. However, upper-case P is widely used. The usage of P vs p depends upon the field in one is working, on the nearby presence of other symbols for quantities such as power and momentum. Mathematically, p = F A where, p is the pressure, F is the normal force and it relates the vector surface element with the normal force acting on it. It is incorrect to say the pressure is directed in such or such direction, the pressure, as a scalar, has no direction. The force given by the relationship to the quantity has a direction. If we change the orientation of the element, the direction of the normal force changes accordingly. Pressure is distributed to solid boundaries or across arbitrary sections of normal to these boundaries or sections at every point. It is a parameter in thermodynamics, and it is conjugate to volume. The SI unit for pressure is the pascal, equal to one newton per square metre and this name for the unit was added in 1971, before that, pressure in SI was expressed simply in newtons per square metre. Other units of pressure, such as pounds per square inch, the CGS unit of pressure is the barye, equal to 1 dyn·cm−2 or 0.1 Pa. Pressure is sometimes expressed in grams-force or kilograms-force per square centimetre, but using the names kilogram, gram, kilogram-force, or gram-force as units of force is expressly forbidden in SI. The technical atmosphere is 1 kgf/cm2, since a system under pressure has potential to perform work on its surroundings, pressure is a measure of potential energy stored per unit volume. It is therefore related to density and may be expressed in units such as joules per cubic metre. Similar pressures are given in kilopascals in most other fields, where the prefix is rarely used

10.
State function
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State functions do not depend on the path by which the system arrived at its present state. A state function describes the state of a system. In contrast, mechanical work and heat are process quantities or path functions, the mode of description breaks down for quantities exhibiting hysteresis effects. A thermodynamic system is described by a number of parameters which are not necessarily independent. The number of parameters needed to describe the system is the dimension of the space of the system. For example, a gas having a fixed number of particles is a simple case of a two-dimensional system. In this example, any system is specified by two parameters, such as pressure and volume, or perhaps pressure and temperature. They are simply different coordinate systems in the thermodynamic state space. Given pressure and temperature, the volume is calculable from them, likewise, given pressure and volume, the temperature is calculable from them. An analogous statement holds for higher-dimensional spaces, as described by the state postulate, if the state space is two-dimensional as in the above example, one may visualize the state space as a three-dimensional graph. The labels of the axes are not generally unique, since there are more variables than three in this case, and any two independent variables suffice to define the state. When a system changes state continuously, it out a path in the state space. The path can be specified by noting the values of the parameters as the system traces out the path, perhaps as a function of time. For example, we might have the pressure P and the volume V as functions of time from time t 0 to t 1 and this will specify a path in our two dimensional state space example. We can now all sorts of functions of time which we may integrate over the path. It is clear that in order to calculate the work W in the integral, we will have to know the functions P and V at each time t. A state function is a function of the parameters of the system which only depends upon the values at the endpoints of the path. For example, suppose we wish to calculate the work plus the integral of V d P over the path, the product P V is therefore a state function of the system

11.
Temperature
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A temperature is an objective comparative measurement of hot or cold. It is measured by a thermometer, several scales and units exist for measuring temperature, the most common being Celsius, Fahrenheit, and, especially in science, Kelvin. Absolute zero is denoted as 0 K on the Kelvin scale, −273.15 °C on the Celsius scale, the kinetic theory offers a valuable but limited account of the behavior of the materials of macroscopic bodies, especially of fluids. Temperature is important in all fields of science including physics, geology, chemistry, atmospheric sciences, medicine. The Celsius scale is used for temperature measurements in most of the world. Because of the 100 degree interval, it is called a centigrade scale.15, the United States commonly uses the Fahrenheit scale, on which water freezes at 32°F and boils at 212°F at sea-level atmospheric pressure. Many scientific measurements use the Kelvin temperature scale, named in honor of the Scottish physicist who first defined it and it is a thermodynamic or absolute temperature scale. Its zero point, 0K, is defined to coincide with the coldest physically-possible temperature and its degrees are defined through thermodynamics. The temperature of zero occurs at 0K = −273. 15°C. For historical reasons, the triple point temperature of water is fixed at 273.16 units of the measurement increment, Temperature is one of the principal quantities in the study of thermodynamics. There is a variety of kinds of temperature scale and 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, 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 capillary tube, is dependent largely on temperature. Such scales are only within convenient ranges of temperature. For example, above the point of mercury, a mercury-in-glass thermometer is impracticable. A material is of no use as a thermometer near one of its phase-change temperatures, 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. Theoretically based temperature scales are based directly on theoretical arguments, especially those of thermodynamics, kinetic theory and they rely on theoretical properties of idealized devices and materials

12.
Thermodynamic free energy
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The thermodynamic free energy is the amount of work that a thermodynamic system can perform. The concept is useful in the thermodynamics of chemical or thermal processes in engineering, the free energy is the internal energy of a system minus the amount of energy that cannot be used to perform work. This unusable energy is given by the entropy of a multiplied by the temperature of the system. Like the internal energy, the energy is a thermodynamic state function. Energy is a generalization of free energy, since energy is the ability to do work which is free energy, free energy is that portion of any first-law energy that is available to perform thermodynamic work, i. e. work mediated by thermal energy. Free energy is subject to loss in the course of such work. Since first-law energy is conserved, it is evident that free energy is an expendable. Several free energy functions may be formulated based on system criteria, free energy functions are Legendre transformations of the internal energy. The Helmholtz free energy has a theoretical importance since it is proportional to the logarithm of the partition function for the canonical ensemble in statistical mechanics. The historically earlier Helmholtz free energy is defined as A = U − TS, where U is the energy, T is the absolute temperature. Its change is equal to the amount of work done on, or obtainable from. Thus its appellation work content, and the designation A from Arbeit, the Gibbs free energy is given by G = H − TS, where H is the enthalpy. Historically, these terms have been used inconsistently. In physics, free energy most often refers to the Helmholtz free energy, denoted by A, while in chemistry, since both fields use both functions, a compromise has been suggested, using A to denote the Helmholtz function and G for the Gibbs function. While A is preferred by IUPAC, G is sometimes still in use, the use of the words “latent heat” implied a similarity to latent heat in the more usual sense, it was regarded as chemically bound to the molecules of the body. In the adiabatic compression of a gas, the heat remained constant. During the early 19th century, the concept of perceptible or free caloric began to be referred to as “free heat” or heat set free. In 1824, for example, the French physicist Sadi Carnot, in his famous “Reflections on the Motive Power of Fire”, an increasing number of books and journal articles do not include the attachment “free”, referring to G as simply Gibbs energy

13.
Thermodynamic temperature
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Thermodynamic temperature is the absolute measure of temperature and is one of the principal parameters of thermodynamics. Thermodynamic temperature is defined by the law of thermodynamics in which the theoretically lowest temperature is the null or zero point. At this point, absolute zero, the constituents of matter have minimal motion. In the quantum-mechanical description, matter at absolute zero is in its ground state, the International System of Units specifies a particular scale for thermodynamic temperature. It uses the Kelvin scale for measurement and selects the point of water at 273.16 K as the fundamental fixing point. Other scales have been in use historically, the Rankine scale, using the degree Fahrenheit as its unit interval, is still in use as part of the English Engineering Units in the United States in some engineering fields. ITS-90 gives a means of estimating the thermodynamic temperature to a very high degree of accuracy. Internal energy is called the heat energy or thermal energy in conditions when no work is done upon the substance by its surroundings. Internal energy may be stored in a number of ways within a substance, each way constituting a degree of freedom. At equilibrium, each degree of freedom will have on average the energy, k B T /2 where k B is the Boltzmann constant. Temperature is a measure of the random submicroscopic motions and vibrations of the constituents of matter. These motions comprise the internal energy of a substance, more specifically, the thermodynamic temperature of any bulk quantity of matter is the measure of the average kinetic energy per classical degree of freedom of its constituent particles. Translational motions are almost always in the classical regime, translational motions are ordinary, whole-body movements in three-dimensional space in which particles move about and exchange energy in collisions. Figure 1 below shows translational motion in gases, Figure 4 below shows translational motion in solids, Zero kinetic energy remains in a substance at absolute zero. Throughout the scientific world where measurements are made in SI units, many engineering fields in the U. S. however, measure thermodynamic temperature using the Rankine scale. By international agreement, the kelvin and its scale are defined by two points, absolute zero, and the triple point of Vienna Standard Mean Ocean Water. Absolute zero, the lowest possible temperature, is defined as being precisely 0 K, the triple point of water is defined as being precisely 273.16 K and 0.01 °C. This definition does three things, It fixes the magnitude of the unit as being precisely 1 part in 273.15 kelvins

14.
Volume (thermodynamics)
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In thermodynamics, the volume of a system is an important extensive parameter for describing its thermodynamic state. The specific volume, a property, is the systems volume per unit of mass. Volume is a function of state and is interdependent with other properties such as pressure and temperature. For example, volume is related to the pressure and temperature of a gas by the ideal gas law. The physical volume of a system may or may not coincide with a control volume used to analyze the system, the volume of a thermodynamic system typically refers to the volume of the working fluid, such as, for example, the fluid within a piston. Changes to this volume may be made through an application of work, an isochoric process however operates at a constant-volume, thus no work can be produced. Many other thermodynamic processes will result in a change in volume, a polytropic process, in particular, causes changes to the system so that the quantity p V n is constant. Note that for specific polytropic indexes a polytropic process will be equivalent to a constant-property process, for instance, for very large values of n approaching infinity, the process becomes constant-volume. Gases are compressible, thus their volumes may be subject to change during thermodynamic processes, liquids, however, are nearly incompressible, thus their volumes can be often taken as constant. In general, compressibility is defined as the volume change of a fluid or solid as a response to a pressure. Similarly, thermal expansion is the tendency of matter to change in volume in response to a change in temperature, many thermodynamic cycles are made up of varying processes, some which maintain a constant volume and some which do not. A vapor-compression refrigeration cycle, for example, follows a sequence where the refrigerant fluid transitions between the liquid and vapor states of matter, typical units for volume are m 3, l, and f t 3. Mechanical work performed on a working fluid causes a change in the constraints of the system, in other words, for work to occur. Hence volume is an important parameter in characterizing many thermodynamic processes where an exchange of energy in the form of work is involved, volume is one of a pair of conjugate variables, the other being pressure. As with all pairs, the product is a form of energy. The product p V is the energy lost to a due to mechanical work. This product is one term which makes up enthalpy H, H = U + p V, the second law of thermodynamics describes constraints on the amount of useful work which can be extracted from a thermodynamic system. Similarly, the value of heat capacity to use in a given process depends on whether the process produces a change in volume