Zeroth law of thermodynamics. The zeroth law of thermodynamics states that if two separate thermodynamic systems are each in thermal equilibrium with a third, then all three are in thermal equilibrium with each other. Two systems are said to be in the relation of thermal equilibrium if they are linked by a wall permeable only to heat, and do not change over time.[1] As a convenience of language, systems are sometimes also said to be in a relation of thermal equilibrium if they are not linked so as to be able to transfer heat to each other, but would not do so if they were connected by a wall permeable only to heat.
The physical meaning of the law was expressed by Maxwell in the words: "All heat is of the same kind".[2] For this reason, another statement of the law is "All diathermal walls are equivalent".[3] The law is important for the mathematical formulation of thermodynamics, which needs the assertion that the relation of thermal equilibrium is an equivalence relation. Zeroth law as equivalence relation[edit] Viscosity. Etymology[edit] The word "viscosity" is derived from the Latin "viscum", meaning "anything sticky, birdlime made from mistletoe, mistletoe".
A viscous glue called birdlime was made from mistletoe berries and was used for lime-twigs to catch birds.[2] Definition[edit] Dynamic (shear) viscosity[edit] Laminar shear of fluid between two plates. In a general parallel flow (such as could occur in a straight pipe), the shear stress is proportional to the gradient of the velocity The dynamic (shear) viscosity of a fluid expresses its resistance to shearing flows, where adjacent layers move parallel to each other with different speeds. . The magnitude of this force is found to be proportional to the speed and the area of each plate, and inversely proportional to their separation The proportionality factor μ in this formula is the viscosity (specifically, the dynamic viscosity) of the fluid.
The ratio where and is the local shear velocity. . , such as in fluid flowing through a pipe. Kinematic viscosity[edit] Third law of thermodynamics. The third law of thermodynamics is sometimes stated as follows, regarding the properties of systems in equilibrium at absolute zero temperature: The entropy of a perfect crystal, at absolute zero kelvin, is exactly equal to zero. The Nernst-Simon statement of the third law of thermodynamics is in regard to thermodynamic processes, and whether it is possible to achieve absolute zero in practice: The entropy change associated with any condensed system undergoing a reversible isothermal process approaches zero as temperature approaches 0 K, where condensed system refers to liquids and solids.
A simpler formulation of the Nernst-Simon statement might be: It is impossible for any process, no matter how idealized, to reduce the entropy of a system to its absolute-zero value in a finite number of operations. Physically, the Nernst-Simon statement implies that it is impossible for any procedure to bring a system to the absolute zero of temperature in a finite number of steps.[3] History[edit] Hence: Thermodynamic system. A thermodynamic system is a precisely specified macroscopic region of the universe, defined by boundaries or walls of particular natures, together with the physical surroundings of that region, which determine processes that are allowed to affect the interior of the region, studied using the principles of thermodynamics. All space in the universe outside the thermodynamic system is known as the surroundings, the environment, or a reservoir.
A system is separated from its surroundings by a boundary, which may be notional or real but, by convention, delimits a finite volume. Transfers of work, heat, or matter and energy between the system and the surroundings may take place across this boundary. A thermodynamic system is classified by the nature of the transfers that are allowed to occur across its boundary, or parts of its boundary. A thermodynamic system has a characteristic set of thermodynamic parameters, or state variables. Overview[edit] History[edit] Boundary[edit] Surroundings[edit] Thermodynamic state. For thermodynamics, a thermodynamic state of a system is fully identified by values of a suitable set of parameters known as state variables, state parameters or thermodynamic variables.
Once such a set of values of thermodynamic variables has been specified for a system, the values of all thermodynamic properties of the system are uniquely determined. Thermodynamics sets up an idealized formalism that can be summarized by a system of postulates of thermodynamics. Thermodynamic states are amongst the fundamental or primitive objects or notions of the formalism, in which their existence is formally postulated, rather than being derived or constructed from other concepts.[1][2] A thermodynamic system is not simply a physical system.[3] Rather, in general, indefinitely many different alternative physical systems comprise a given thermodynamic system, because in general a physical system has vastly many more detailed characteristics than are mentioned in a thermodynamic description.
Thermodynamic process. A thermodynamic process is the energetic development of a thermodynamic system proceeding from an initial state to a final state. Paths through the space of thermodynamic variables are often specified by holding certain thermodynamic variables constant. A state function is a thermodynamic variable which depends only on the current state of the system, not the path taken to reach that state. Conversely a process function does depend on the path. Overview[edit] An example of a series of thermodynamic processes which make up the Stirling cycle Conjugate variable processes[edit] Pressure - volume[edit] The pressure-volume conjugate pair is concerned with the transfer of mechanical or dynamic energy as the result of work.
An isobaric process occurs at constant pressure. Temperature - entropy[edit] The temperature-entropy conjugate pair is concerned with the transfer of thermal energy as the result of heating. An isothermal process occurs at a constant temperature. Thermodynamic potentials[edit] Thermodynamic potential. Description and interpretation[edit] Five common thermodynamic potentials are:[1] These five common potentials are all energy potentials, but there are also entropy potentials.
The thermodynamic square can be used as a tool to recall and derive some of the potentials. Just as in mechanics, where potential energy is defined as capacity to do work, similarly different potentials have different meanings. In particular: (see principle of minimum energy for a derivation)[3] When the entropy (S ) and "external parameters" (e.g. volume) of a closed system are held constant, the internal energy (U ) decreases and reaches a minimum value at equilibrium. Natural variables[edit] If there is only one species, then we are done. And so on. The fundamental equations[edit] where δQ is the infinitesimal heat flow into the system, and δW is the infinitesimal work done by the system, μi is the chemical potential of particle type i and Ni is the number of type i particles. Where T is temperature, S is entropy, Thermodynamic equilibrium. In thermodynamics, a thermodynamic system is in thermodynamic equilibrium when it is in thermal equilibrium, mechanical equilibrium, radiative equilibrium, and chemical equilibrium.
Equilibrium means a state of balance. In a state of thermodynamic equilibrium, there are no net flows of matter or of energy, no phase changes, and no unbalanced potentials (or driving forces), within the system. A system that is in thermodynamic equilibrium experiences no changes when it is isolated from its surroundings. In non-equilibrium systems there are net flows of matter or energy, or phase changes are occurring; if such changes can be triggered to occur in a system in which they are not already occurring, it is said to be in a metastable equilibrium. When a body of material starts from a non-equilibrium state of inhomogeneity or chemical non-equilibrium, and is then isolated, it spontaneously evolves towards its own internal state of thermodynamic equilibrium.
Overview[edit] Reservations[edit] J.A. R. Temperature. A map of global long term monthly average surface air temperatures in Mollweide projection. A temperature is a numerical measure of hot and cold. Its measurement is by detection of heat radiation or particle velocity or kinetic energy, or by the bulk behavior of a thermometric material. It may be calibrated in any of various temperature scales, Celsius, Fahrenheit, Kelvin, etc. The fundamental physical definition of temperature is provided by thermodynamics. Measurements with a small thermometer, or by detection of heat radiation, can show that the temperature of a body of material can vary from time to time and from place to place within it.
Within a body that exchanges no energy or matter with its surroundings, temperature tends to become spatially uniform as time passes. The kinetic theory offers a valuable but limited account of the behavior of the materials of macroscopic systems. Thermal vibration of a segment of proteinalpha helix. Use in science[edit] Temperature scales[edit] Statistical ensemble (mathematical physics) In mathematical physics, especially as introduced into statistical mechanics and thermodynamics by J. Willard Gibbs in 1902, an ensemble (also statistical ensemble) is an idealization consisting of a large number of virtual copies (sometimes infinitely many) of a system, considered all at once, each of which represents a possible state that the real system might be in.
In other words, a statistical ensemble is a probability distribution for the state of the system.[1] A thermodynamic ensemble is a specific variety of statistical ensemble that, among other properties, is in statistical equilibrium (defined below), and is used to derive the properties of thermodynamic systems from the laws of classical or quantum mechanics.[2][3] This article treats the notion of ensembles in a mathematically rigorous fashion, although relevant physical aspects will be mentioned. The concept of an equilibrium or stationary ensemble is crucial to some applications of statistical ensembles.
Where. State function. In contrast, mechanical work and heat are process quantities because their values depend on the specific transition (or path) between two equilibrium states. History[edit] It is likely that the term “functions of state” was used in a loose sense during the 1850s and 60s by those such as Rudolf Clausius, William Rankine, Peter Tait, William Thomson, and it is clear that by the 1870s the term had acquired a use of its own. In 1873, for example, Willard Gibbs, in his paper “Graphical Methods in the Thermodynamics of Fluids”, states: “The quantities V, B, T, U, and S are determined when the state of the body is given, and it may be permitted to call them functions of the state of the body.” Overview[edit] ).
For example, a monatomic gas with a fixed number of particles is a simple case of a two-dimensional system ( ). When a system changes state continuously, it traces out a "path" in the state space. And the volume as functions of time from time to to time we calculate and at each time . Spontaneous process. A spontaneous process is the time-evolution of a system in which it releases free energy (usually as heat) and moves to a lower, more thermodynamically stable energy state.[1][2] The sign convention of changes in free energy follows the general convention for thermodynamic measurements, in which a release of free energy from the system corresponds to a negative change in free energy, but a positive change for the surroundings.
Depending on the nature of the process, the free energy is determined differently. For example, the Gibbs free energy is used when considering processes that occur under constant pressure and temperature conditions whereas the Helmholtz free energy is used when considering processes that occur under constant volume and temperature conditions. A spontaneous process is capable of proceeding in a given direction, as written or described, without needing to be driven by an outside source of energy. Overview[edit] See also[edit] References[edit] Second law of thermodynamics. The second law of thermodynamics states that every process occurring in nature proceeds in the sense in which the sum of the entropies of all bodies taking part in the process is increased.
In the limit, i.e. for reversible processes, the sum of the entropies remains unchanged.[1][2][3] The second law is an empirical finding that has been accepted as an axiom of thermodynamic theory. Statistical thermodynamics, classical or quantum, explains the law. The second law has been expressed in many ways. Its first formulation is credited to the French scientist Sadi Carnot in 1824 (see Timeline of thermodynamics). Introduction[edit] The first law of thermodynamics provides the basic definition of thermodynamic energy, also called internal energy, associated with all thermodynamic systems, but unknown in classical mechanics, and states the rule of conservation of energy in nature.[4][5] For mathematical analysis of processes, entropy is introduced as follows.
Various statements of the law[edit] Thus, Reversible process (thermodynamics) In thermodynamics, a reversible process -- or reversible cycle if the process is cyclic -- is a process that can be "reversed" by means of infinitesimal changes in some property of the system without entropy production (i.e. dissipation of energy).[1] Due to these infinitesimal changes, the system is in thermodynamic equilibrium throughout the entire process. Since it would take an infinite amount of time for the reversible process to finish, perfectly reversible processes are impossible.
However, if the system undergoing the changes responds much faster than the applied change, the deviation from reversibility may be negligible. In a reversible cycle, the system and its surroundings will be exactly the same after each cycle.[2] In thermodynamics, processes can be carried out in one of two ways: reversibly or irreversibly. Reversibility in thermodynamics refers to performing a reaction continuously at equilibrium.
Pressure. Pressure as exerted by particle collisions inside a closed container. 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 (also spelled gage pressure)[a][not in citation given] is the pressure relative to the local atmospheric or ambient pressure. Definition[edit] Pressure is the amount of force acting per unit area. The symbol of pressure is p or P. [b][1] Formula[edit] Mathematically: where: is the pressure, is the normal force, is the area of the surface on contact. The minus sign comes from the fact that the force is considered towards the surface element, while the normal vector points outward.
It is incorrect (although rather usual) to say "the pressure is directed in such or such direction". Units[edit] Since a system under pressure has potential to perform work on its surroundings, pressure is a measure of potential energy stored per unit volume. Examples[edit] Scalar nature[edit] via where. Partition function (statistical mechanics) Non-equilibrium thermodynamics. Internal energy. Ideal gas law. Heat. Thermodynamic free energy. First law of thermodynamics. Equipartition theorem. Equation of state. Entropy. Enthalpy. Conjugate variables (thermodynamics) Boltzmann constant.