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Thermodynamics
Thermodynamics
Thermodynamics
is a branch of physics concerned with heat and temperature and their relation to other forms of energy and work. The behavior of these quantities is governed by the four laws of thermodynamics, irrespective of the composition or specific properties of the material or system in question. The laws of thermodynamics are explained in terms of microscopic constituents by statistical mechanics
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Carnot Heat Engine
A Carnot heat engine[2] is a theoretical engine that operates on the reversible Carnot cycle. The basic model for this engine was developed by Nicolas Léonard Sadi Carnot
Nicolas Léonard Sadi Carnot
in 1824. The Carnot engine model was graphically expanded upon by Benoît Paul Émile Clapeyron in 1834 and mathematically explored by Rudolf Clausius
Rudolf Clausius
in 1857 from which the concept of entropy emerged. Every thermodynamic system exists in a particular state. A thermodynamic cycle occurs when a system is taken through a series of different states, and finally returned to its initial state. In the process of going through this cycle, the system may perform work on its surroundings, thereby acting as a heat engine. A heat engine acts by transferring energy from a warm region to a cool region of space and, in the process, converting some of that energy to mechanical work. The cycle may also be reversed
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State Of Matter
In physics, a state of matter is one of the distinct forms in which matter can exist. Four states of matter are observable in everyday life: solid, liquid, gas, and plasma. Many other states are known to exist, such as glass or liquid crystal, and some only exist under extreme conditions, such as Bose–Einstein condensates, neutron-degenerate matter, and quark-gluon plasma, which only occur, respectively, in situations of extreme cold, extreme density, and extremely high-energy. Some other states are believed to be possible but remain theoretical for now. For a complete list of all exotic states of matter, see the list of states of matter. Historically, the distinction is made based on qualitative differences in properties. Matter
Matter
in the solid state maintains a fixed volume and shape, with component particles (atoms, molecules or ions) close together and fixed into place
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List Of Thermodynamic Properties
Within thermodynamics, a physical property is any property that is measurable, and whose value describes a state of a physical system. Some constants, such as the ideal gas constant, R, do not describe the state of a system, and so are not properties. On the other hand, some constants, such as Kf (the freezing point depression constant, or cryoscopic constant), depend on the identity of a substance, and so may be considered to describe the state of a system, and so may be considered physical properties. "Specific" properties are expressed on a per mass basis. If the units were changed from per mass to, for example, per mole, the property would remain as it was (i.e., intensive or extensive). Regarding Work and Heat Work and heat are not thermodynamic properties, but rather process quantities: flows of energy across a system boundary
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Real Gas
Real gases are non-hypothetical gases whose molecules occupy space and have interactions; consequently, they adhere to gas laws. To understand the behaviour of real gases, the following must be taken into account:compressibility effects; variable specific heat capacity; van der Waals forces; non-equilibrium thermodynamic effects; issues with molecular dissociation and elementary reactions with variable compositionFor most applications, such a detailed analysis is unnecessary, and the ideal gas approximation can be used with reasonable accuracy. On the other hand, real-gas models have to be used near the condensation point of gases, near critical points, at very high pressures, to explain the Joule–Thomson effect
Joule–Thomson effect
and in other less usual cases
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Quantum Thermodynamics
Quantum thermodynamics is the study of the relations between two independent physical theories: thermodynamics and quantum mechanics. The two independent theories address the physical phenomena of light and matter. In 1905 Einstein
Einstein
argued that the requirement of consistency between thermodynamics and electromagnetism[1] leads to the conclusion that light is quantized obtaining the relation E = h ν displaystyle E=hnu . This paper is the dawn of quantum theory. In a few decades quantum theory became established with an independent set of rules.[2] Currently quantum thermodynamics addresses the emergence of thermodynamic laws from quantum mechanics. It differs from quantum statistical mechanics in the emphasis on dynamical processes out of equilibrium
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Thermodynamic Diagrams
Thermodynamic
Thermodynamic
diagrams are diagrams used to represent the thermodynamic states of a material (typically fluid) and the consequences of manipulating this material. For instance, a temperature–entropy diagram (T–s diagram) may be used to demonstrate the behavior of a fluid as it is changed by a compressor.Contents1 Overview 2 Types of thermodynamic diagrams 3 Characteristics 4 References 5 Further reading 6 External linksOverview[edit] Especially in meteorology they are used to analyze the actual state of the atmosphere derived from the measurements of radiosondes, usually obtained with weather balloons. In such diagrams, temperature and humidity values (represented by the dew point) are displayed with respect to pressure. Thus the diagram gives at a first glance the actual atmospheric stratification and vertical water vapor distribution
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Isochoric Process
An isochoric process, also called a constant-volume process, an isovolumetric process, or an isometric process, is a thermodynamic process during which the volume of the closed system undergoing such a process remains constant. An isochoric process is exemplified by the heating or the cooling of the contents of a sealed, inelastic container: The thermodynamic process is the addition or removal of heat; the isolation of the contents of the container establishes the closed system; and the inability of the container to deform imposes the constant-volume condition
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Carnot Cycle
The Carnot cycle
Carnot cycle
is a theoretical thermodynamic cycle proposed by French physicist Sadi Carnot in 1824 and expanded upon by others in the 1830s and 1840s. It provides an upper limit on the efficiency that any classical thermodynamic engine can achieve during the conversion of heat into work, or conversely, the efficiency of a refrigeration system in creating a temperature difference (e.g. refrigeration) by the application of work to the system. It is not an actual thermodynamic cycle but is a theoretical construct. Every single thermodynamic system exists in a particular state. When a system is taken through a series of different states and finally returned to its initial state, a thermodynamic cycle is said to have occurred. In the process of going through this cycle, the system may perform work on its surroundings, for example by moving a piston, thereby acting as a heat engine
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Intensive And Extensive Properties
Physical properties of materials and systems can often be categorized as being either intensive or extensive quantities, according to how the property changes when the size (or extent) of the system changes. According to IUPAC, an intensive property is one whose magnitude is independent of the size of the system. An extensive property is one whose magnitude is additive for subsystems.[1] An intensive property is a bulk property, meaning that it is a physical property of a system that does not depend on the system size or the amount of material in the system. Examples of intensive properties include temperature, T, refractive index, n, density, ρ, and hardness of an object, η (IUPAC symbols[1] are used throughout this article)
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Pressure
Pressure
Pressure
(symbol: p or P) is the force applied perpendicular to the surface of an object per unit area over which that force is distributed. Gauge pressure
Gauge pressure
(also spelled gage pressure)[a] is the pressure relative to the ambient pressure. Various units are used to express pressure. Some of these derive from a unit of force divided by a unit of area; the SI unit
SI unit
of pressure, the pascal (Pa), for example, is one newton per square metre; similarly, the pound-force per square inch (psi) is the traditional unit of pressure in the imperial and US customary systems. Pressure may also be expressed in terms of standard atmospheric pressure; the atmosphere (atm) is equal to this pressure, and the torr is defined as ​1⁄760 of this
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Quasistatic Process
In thermodynamics, a quasi-static process is a thermodynamic process that happens slowly enough for the system to remain in internal equilibrium. An example of this is quasi-static compression, where the volume of a system changes at a rate slow enough to allow the pressure to remain uniform and constant throughout the system.[1] Any reversible process is necessarily a quasi-static one. However, quasi-static processes involving entropy production are irreversible. An example of a quasi-static process that is not reversible is a compression against a system with a piston subject to friction—although the system is always in thermal equilibrium, the friction ensures the generation of dissipative entropy, which directly goes against the definition of reversible
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Polytropic Process
A polytropic process is a thermodynamic process that obeys the relation: p V n = C displaystyle pV^ ,n =C where p is the pressure, V is volume, n is the polytropic index , and C is a constant
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Conjugate Variables (thermodynamics)
In thermodynamics, the internal energy of a system is expressed in terms of pairs of conjugate variables such as temperature and entropy or pressure and volume. In fact, all thermodynamic potentials are expressed in terms of conjugate pairs. The product of two quantities that are conjugate has units of energy or sometimes power. For a mechanical system, a small increment of energy is the product of a force times a small displacement. A similar situation exists in thermodynamics. An increment in the energy of a thermodynamic system can be expressed as the sum of the products of certain generalized "forces" that, when unbalanced, cause certain generalized "displacements", and the product of the two is the energy transferred as a result. These forces and their associated displacements are called conjugate variables. The thermodynamic force is always an intensive variable and the displacement is always an extensive variable, yielding an extensive energy transfer
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Free Expansion
Free expansion
Free expansion
is an irreversible process in which a gas expands into an insulated evacuated chamber. It is also called Joule expansion. Real gases experience a temperature change during free expansion. For an ideal gas, the temperature doesn't change, and the conditions before and after adiabatic free expansion satisfy ( P i ) ( V i ) = ( P f ) ( V f ) displaystyle (P_ i )(V_ i )=(P_ f )(V_ f ) , where p is the pressure, V is the volume, and i and f refer to the initial and final states. Since the gas expands, Vf > Vi , which implies that the pressure does drop (Pf < Pi). During free expansion, no work is done by the gas
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Isobaric Process
An isobaric process is a thermodynamic process in which the pressure stays constant: ΔP = 0. The heat transferred to the system does work, but also changes the internal energy of the system. This article uses the chemistry sign convention for work, where positive work is work done on the system. Using this convention, by the first law of thermodynamics,The yellow area represents the work done Q = Δ U − W displaystyle Q=Delta U-W, where W is work, U is internal energy, and Q is heat.[1] Pressure-volume work by the closed system is defined as: W = − ∫ p d V displaystyle W=-int !p,dV, where Δ means change over the whole process, whereas d denotes a differential
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