Loading AI tools
User:Chjoaygame/sandbox/archive 1
This is the user sandbox of Chjoaygame. A user sandbox is a subpage of the user's user page. It serves as a testing spot and page development space for the user and is not an encyclopedia article. Create or edit your own sandbox here. Other sandboxes: Main sandbox | Template sandbox Finished writing a draft article? Are you ready to request review of it by an experienced editor for possible inclusion in Wikipedia? Submit your draft for review! |
Xxx〈bra|ket〉 xxX
|p⟩
text in blue font
ν = c/λ
Gibb's mixing paradox concerns mixing of two distinct chemical species that scarcely differ, in comparison with what happens when two bodies of identical material are put in direct contact without a separating wall. In the case of the distinct species, a significantly large entropy change may occur. But none in the case of identical material. The question arises, of what happens as the difference of species decreases continuously to zero. One physical answer to this question is that no such continuous decrease to zero is possible, because chemical species are defined quantally. That is the quantum theory, and is generally accepted as correct.
The second law of thermodynamics states that when a thermodynamically defined process of transfers of matter and energy occurs, the sum of the entropies of the participating bodies must increase. In an idealized limiting case, that of a reversible process, this sum remains unchanged.
In a thermodynamically defined process of transfers between bodies of matter and radiation, each participating body is initially in its own state of internal thermodynamic equilibrium. The bodies are initially separated from one another by walls that obstruct the passage of matter and energy between them. The transfers are initiated by a thermodynamic operation: some external agency makes one or more of the walls less obstructive.[1] This establishes new equilibrium states in the bodies. If, instead of making the walls less obstructive, the thermodynamic operation makes them more obstructive, there is no effect on an established thermodynamic equilibrium.
The law expresses the irreversibility of the process. The transfers invariably bring about spread,[2][3][4] dispersal, or dissipation[5] of matter or energy, or both, amongst the bodies. They occur because more kinds of transfer through the walls have become possible.[6] Irreversibility in thermodynamic processes is a consequence of the asymmetric character of thermodynamic operations, and not of any internally irreversible microscopic properties of the bodies. Thermodynamic operations are macroscopic external interventions imposed on the participating bodies, not derived from their internal properties.
The second law is an empirical finding that has been accepted as an axiom of thermodynamic theory. Of course, thermodynamics relies on presuppositions, for example the existence of perfect thermodynamic equilibrium, that are not exactly fulfilled by nature. Statistical thermodynamics, classical or quantum, explains the microscopic origin of 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).
To comply with Editor [[User:|User:]]'s concern about convoluted wording quoted from a translation of Planck, I am posting a re-worded version of the Planck statement, as follows: "The second law of thermodynamics states that in every natural thermodynamic process the sum of the entropies of all participating bodies is increased. In the limiting case, for reversible processes this sum remains unchanged."
There are several reasons why Planck's statement of the law is excellent.
There are some background considerations, as follows.
Here Editor Ppithermo, with concurrence here from Editor PAR, considers an isolated system with an internal adiabatic wall, for example a piston that allows transfer as work, but not of energy as heat, or of matter, that creates internal adiabatic enclosures. He refers to problem 2.7–3 in the second edition of Callen.[1] He had brought up the discussion of this problem in the first edition of Callen, in Appendix C, headed Equilibrium with internal adiabatic constraints.[2]
Callen further indicates this point by his use of the proviso "in the absence of an internal constraint" in his statement of his
Bailyn is less explicit, but observes the same point in what he calls the "superstrong entropy" version of the law:
Landsberg[5] discusses a distinction between "simple" and "compound systems". Compound systems contain, for example, simple systems separated by adiabatic walls. Landsberg writes about equilibria of compound systems: "... This can occur if certain partitions separate the systems, or certain constraints are part of the arrangement, which are capable of preventing a thermodynamic process which would otherwise be possible."[6] Landsberg sets a student exercise: "5.11. Decide the following finer points of thermodynamics, giving reasons: ... (b) How good a statement is 'The entropy of an isolated system cannot decrease with time'?[7]
These items are why it is not part of the usual statements of the second law to claim that "... isolated systems always evolve toward thermodynamic equilibrium, the state with maximum entropy." True, those just quoted words are rather vague, and they can be pressed to read so as to make them not strictly false. But they have a strong potential to mislead, and express an idea that is not explicit in the standard statements of the second law. They should not be in the first sentence of the lead, that is more or less a definition.
In the background of the Planck version, there is the following thinking.
The process starts with a state of thermodynamic equilibrium between several systems, each obeying the minus oneth law of thermodynamics, separated by walls of specific permeabilities. Then a thermodynamic operation occurs, marking the transition from initial to final states, changing the permeabilities, and allowing initially prohibited transfers. There follows a final state of thermodynamic equilibrium. Entropy sums are defined and compared for the initial and final states.
The Planck version avoids talk about some idea of "progressive advance towards equilibrium". This is desirable because an equilibrium state is of practically infinite duration and is subject to recurrence. A transient "initial disequilibrium" can be ignored for a state of infinite duration. Entropy is defined for states of thermodynamic equilibrium.
The second law is conveniently stated in terms of a quantity called 'entropy'. The present section of this article intends to explain how this is so.
The concept of entropy rests on some presuppositions:
One may consider two isolating enclosures separated by an impermeable wall. Initially, one contains a gas in its own internal state of thermodynamic equilibrium, the other a vacuum. The entropy of the one may be denoted S1. From the scaling and additive properties, the other has entropy 0. The entropy of the initially compartmented compound system, from the additive property, is S+ = S1 + 0 = S1. Now a thermodynamic operation may remove the partition. In the consequent new equilibrium, according to the second law, the entropy Su of the new unpartitioned system obeys Su > S+ = S1.
By a second thermodynamic operation, the partition may then be replaced. By the scaling property, the entropy of each new half-system is S1/2 = Su/2. The entropy of the compound of the two separated half-systems is Sc = 2S1/2 = Su, unchanged by the operation.
The original two-compartment compound system can be restored by adiabatic compression by a force in the surroundings, acting on the expanded compound system, followed by heat transfer to or from the surroundings. A force in the surroundings can compress the gas only in accord with the second law, increasing the entropy sum of the compound system and the surroundings. Also the work transfer increases the internal energy of the comppund system. It is conventional to suppose that the adiabatic work is suppplied by a source that does not thereby increase its own entropy. Then the entropy of the adiabatically compressed system, before heat transfer, has an entropy that satisfies Sa > Sc = Su > S1. It is as if the work transfer was associated with the "creation" of some entropy inside the compound system. The transfer of heat must by the first law be from the compound system to the surroundings, so as to restore the original internal energy. It is as if the heat transfer is accompanied by a "transfer of entropy", as if entropy were some kind of material, that could be transferred; rigorously considered, this does not really make sense, but it is, for the present, a harmless metaphor.
"The starting point is undoubtedly the qualitative connection between the temperature concept and our crude physiological sensations of hot and cold."[1]
Wall (Impermeable) | External relation | ||||||||||
Rigid | T-Ins | No IW | |||||||||
δQ | δWx | dT | dP | dS | dV | dN | dV=0 | δQ=0 | δWx=0 | ||
Process | |||||||||||
Isentropic process | 0 | 0 | - | - | 0 | - | 0 | no | yes | yes | closed system |
Adiabatic process | 0 | - | - | - | - | - | 0 | no | yes | no | closed system |
Adynamic process | - | 0 | - | - | - | 0 | 0 | yes | no | yes | closed system |
Isochoric process | - | - | - | - | - | 0 | - | yes | - | - | constant volume system |
Isobaric process | - | - | - | 0 | - | - | - | - | - | - | constant pressure reservoir |
Isothermal process | - | - | 0 | - | - | - | - | - | no | - | constant temperature reservoir |
x-Isolated System | |||||||||||
Adiabatic Is | 0 | - | - | - | - | - | 0 | - | yes | - | closed system |
Closed system | - | - | - | - | - | - | 0 | - | - | - | closed system |
Isolated system | 0 | 0 | - | 0 | - | 0 | 0 | yes | yes | yes | isolated system |
Open system | - | - | - | - | - | - | - | no | no | no | open system |
Conserved Thermodynamic potential | |||||||||||
U: Internal energy | 0 | 0 | - | 0 | - | 0 | 0 | yes | yes | yes | - |
F: Helmholtz free energy | - | - | 0 | - | - | 0 | 0 | yes | no | - | constant temperature reservoir |
H: Enthalpy | 0 | 0 | - | 0 | 0 | - | 0 | no | yes | yes | constant pressure reservoir |
G: Gibbs free energy | - | - | 0 | - | 0 | - | 0 | no | no | - | constant pressure and temperature reservoir |
δQ is heat, δWx is irreversible work, so TdS=δQ+δWx. If both are zero, then dS=0. For the 3 possible walls, "T ins" means thermally insulated, and "No IW" means no irreversible work. "External intensive variable" specifies the region external to the system.
Einstein (1936, Physics and Reality:
This is very close to Whitehead's ontological principle. Translation touched up by me.
Agrees with Whitehead.
Established in 1875, the London and Newcastle Tea Company's offices were in Charlotte Square, Newcastle. Early branches were in New Bridge Street, Sandhill, Scotswood Road, Shields Road, Westgate Road and Clayton Street but the firm later expanded to more outlying parts of the city. It went out of business between 1959 and 1962.[1]
According to James B. Jefferys, the London and Newcastle Tea Company was one of the first five firms in Britain in the grocery and provisions trade to operate in a multiple branch scale. In 1875, it ran between 10 and 20 branches; in 1880, it was the second biggest grocery chain in Britain, with between 40 and 50 branches, just ahead of the rapidly expanding Thomas Lipton.[2]
The firm had a loyalty scheme in operation as early as 1875. The network of groceries, which sold the company’s tea, gave a brass check with each purchase. Customers were invited to save the checks until they had acquired enough to claim a prize such as a toy, an item of crockery or a household gadget. The checks are now collectors’ items. By 1928, the shop at 212 Chillingham Road, Heaton, had been acquired by the London and Newcastle Tea Company.[3]
Guggenheim writes
E.T. Jaynes has expressed opinion on this topic.
Also
Denbigh writes
Anderson writes
Atkins & de Paula write
Bridgman writes
Adkins writes
Attard writes
Baierlein writes
Brush writes
Callen writes
Multiple causes
Any event may have several causes. This possibility is not excluded by my definition (given explicitly on p. 9), though I speak there of A being 'the' cause of the effect B. Actually the 'number' of causes, i.e. of conditions on which an effect B depends, seems to me a rather meaningless notion. One often finds the idea of a 'causal chain' A1, A2, ••• , where B depends directly on A1, A1 on A2, etc., so that B depends indirectly on any of the An. As the series may never end — where is a 'first cause' to be found ? — the number of causes may be, and will be in general, infinite. But there seems to be not the slightest reason to assume only one such chain, or even a number of chains; for the causes may be interlocked in a complicated way, and a 'network' of causes (even in a multi-dimensional space} seems to be a more appropriate picture. Born, Max (1949), Natural Philosophy of Cause and Chance, Oxford University Press, London UK, page 129.
Editor Incnis Mrsi has made an edit here.
The present article is about the thermodynamic concept of heat transfer, which is carefully defined in thermodynamics. This should be made clear in the article, which for a general readership should start from elementary considerations. So defined, heat transfer by convection must be through a process with a mechanism that results in no net transfer of matter: the convection must be circulatory. Though implicit in them, this is not explicitly emphasized in the articles Convective heat transfer and Natural convection. Also important here is that convective circulation is spontaneous only beyond some suitably specified threshold.
The thermodynamic definition of heat refers to a process of transfer of energy that causes a change of a body from one state of its own internal thermodynamic equilibrium to another. In the early days of thermodynamics, this was expressed in terms of cyclic processes. The notion of thermodynamic equilibrium requires zero macroscopic flows. The articles about Convective heat transfer and about Natural convection are mainly about flows, and are specialized developments beyond elementary thermodynamics.
A thermodynamic process is initiated and terminated by thermodynamic operations. In the present context, one may envisage three bodies, a source body, a fluid intermediary, and a destination body. The heat is transferred from source to destination, eventually making no change to the fluid intermediary. This may be arranged through thermodynamic operations by which the walls separating the three bodies are changed from adiabatic to diathermal (initiating) and back from diathermal to adiabatic (terminating). The source body starts at a higher temperature than does the destination body. Such a temperature difference is conceptually proper to elementary thermodynamics, which does not deal with the dynamics of the flows by which the heat is transferred. For the thermodynamic definition of heat transfer, such a temperature difference is suitable as a threshold parameter. Temperature gradients entail flow, and do not belong to the thermodynamic definition of heat transfer.
The edit by Incnis Mrsi is concerned with dynamic flow variables internal to the fluid intermediary of the process of convective circulation, not with the respective initial and final temperatures of the source and destination bodies. It is conceptually specialized beyond an elementary or fundamental approach to thermodynamics. The edit belongs to topics such as fluid dynamics, which deal with the dynamic stability of flows, beyond the scope of the four laws of thermodynamics. For example, chapter 2 of Hydrodynamic and Hydromagnetic Stability, the monumental 1961 textbook by Chandrasekhar,[1] treats the onset of convective circulation in a special case. That chapter mentions neither of the touchstones of thermodynamics, entropy and the second law of thermodynamics, because it is not about thermodynamics, the domain of the present article.
For a Wikipedia article, it is valuable to present material with clarity of principle, because Wikipedia is partly pedagogic. Consequently, it is worth distinguishing between the respective scopes of thermodynamics and fluid mechanics because of Einstein's dictum that, within its domain of applicability, thermodynamics is the one area of physics that will never be overthrown. A factor of the permanent reliability of thermodynamics is that its domain of applicability is strictly defined and confined.
Placed in the lead of the present article, the edit will confuse the thinking of ordinary readers.
The previous entry was "Though spontaneous, convective circulation does not necessarily and immediately occur merely because of temperature difference; for it to occur in a given arrangement of systems, there is a threshold temperature difference that needs to be exceeded."
The new edit reads "Natural convection, though spontaneous, does not occur merely because of temperature difference but due to Rayleigh instability caused by combination of gravity and thermal expansion in supplement to sufficient temperature gradient."
Non-convective transfer of energy as heat has two kinds: passive, as in, for example, conduction and radiation; and active, as in, for example, Joule's original mechanical-shaft-and-paddle experiment, magnetic stirring, and Joule heating.
In passive non-convective heat transfer between two thermodynamic systems, source and destination transfers are of the same character. Each thermodynamic system, starting from an initial state of homogeneous internal thermodynamic equilibrium, can export or import energy as heat in this way. Transfer is governed merely by temperature difference between source and destination bodies, and is spontaneous, immediate in onset, and necessary, whenever there is a suitable pathway for transfer.
In active non-convective heat transfer between two thermodynamic systems, source and destination transfers are of different characters. The source energy transfer is as work done by forces originating in energy-exporting system, and the destination energy transfer is as heating of the energy-importing system. A thermodynamic system, starting from an initial state of internal thermodynamic equilibrium, cannot export energy as heat in this way. Active non-convective heat transfer is not governed by temperature difference between source and destination bodies. For active non-convective heat transfer, the source system must be in an inhomogeneous internal equilibrium. A system in a homogeneous internal equilibrium cannot actively export energy as heat.
Neither of these two kinds of non-convective heat transfer has a temperature threshold that needs to be exceeded.
Also, convective circulatory heat transfer has two kinds: spontaneous, and driven.
As defined in thermodynamics, spontaneous circulatory convective heat transfer has a feature different from non-convective heat transfer: in a given arrangement of systems and surroundings, it occurs only when a threshold difference in the temperatures of the source and destination bodies is exceeded.
In driven heat spreading by circulatory convection, the driving work is done by agencies in the intermediary surroundings of the source and destination bodies, at least one of which should exhibit passive heat transfer. It may be a fine point to require that passive heat transfer occur at both source and destination. One could allow that, at just one of source or destination, energy transfer between convective fluid and body of interest involve transfers of matter that preclude a proper and unique definition of transfer as heat. For an example in nature, in the transfer of energy from working muscles to the surroundings through the skin, carried by the blood stream, driven by the work of the heart, the transfer from skin to surroundings by radiation is passive heat transfer, while the transfer of energy from muscles to blood will in general involve matter transfers so that heat transfer cannot be uniquely separated or defined. Unlike the spontaneous kind of circulatory convective heat transfer, driven circulatory convective heat transfer is not necessarily subject to a threshold difference in the temperatures of the source and destination bodies.
The thermodynamic definition of heat transfer is by exclusion of thermodynamic work and matter transfer. Quantitatively, it is expressed in terms of thermodynamic variables such as temperature and entropy, which are defined for source and destination bodies initially and finally in their own states of internal thermodynamic equilibrium, which by definition requires zero macroscopic flows. For spontaneous circulatory convection in a given arrangement of systems, calculations of the temperature difference threshold are not directly based on the laws of thermodynamics . Such calculations require non-thermodynamic variables, in particular, those of fluid dynamics, beyond the scope of this present article.
An edit has replaced the entry
with the entry
The edit cover note reads “because it is not pressure–volume work” was an exemplary poor logic, wasn’t it?.
The logic of the former entry seems fine to me. What is wrong with it? The only kind of thermodynamic work that is relevant here is pressure–volume work. The relevant externally applied macroscopic mechanical work is not pressure–volume work. Therefore it is not the kind of thermodynamic work that is relevant here.
The new phrase "non-equilibrium process" is likely to confuse or even mislead the ordinary reader of Wikipedia. At a glance, the new phrase may seem to be a customary term of art in thermodynamics, but it is not so. A thermodynamic process is a change in a thermodynamic system from one state of thermodynamic equilibrium to another. Consequently, every thermodynamic process involves a departure from, and an entry into, some thermodynamic equilibrium or other, and in that sense might be called 'non-equilibrium', so that it would make no sense to speak of an 'equilibrium process'; then the phrase 'non-equilibrium process' appears as an uninformative tautology, unhelpful for the reader. The term 'reversible process' is used often enough, and perhaps the term "non-equilibrium process" is intended refer to a process that is not reversible. The term 'reversible process', however, is itself suggestive of misconception, unless it is accompanied by a reminder that every actually occurring thermodynamic process is irreversible, and that the term refers to a fictive, conceptual, or mathematical idealization. In short, the phrase "non-equilibrium process" is unfit for the present Wikipedia purpose.
In context, pressure and volume are the relevant state variables or state functions. That the work "cannot be described with state variables" evidently means just that it is not pressure–volume work.
The new entry is not as suitable as the one it replaces, and may be misleading.
This book uses the word 'matter' as ordinary language, but, in near, though not exact, agreement with the poster of the edit, it once uses the phrase 'transfer of mass'. It writes neither 'mass transfer', nor 'transfer of matter'.
Page 222:
Page 30:
Page 121:
Page 131:
Page 157:
Page 270:
This book uses the word 'matter' as ordinary language. It uses the phrases 'matter transfer' and 'transfer of matter', but does not use the phrase 'mass transfer'.
Page 5:
Page 17:
Page 26:
Page 58:
This book uses the word 'matter' as ordinary language. It uses the phrases 'transfer of matter', and 'transfer of material', but not the phrase 'mass transfer'.
Page 5:
Page 6:
Page 77:
Page 81:
This book does not use the phrase 'mass transfer' where the poster of the edits would expect it. This book uses the word 'matter' as in ordinary language, and the phrase 'flow of matter' in a place where the poster of the edit would expect 'mass transfer'.
Page 51:
Page 59:
Page 157:
Page 327:
Page 328:
Page 338:
Page 341:
Page 342:
Page 343:
Page 345:
Page 348:
Page 366:
This book does not use the phrase 'mass transfer' where the poster of the edits would expect it. This book uses the word 'matter' as in ordinary language, and the phrases or clauses 'passage of matter', 'exchange of matter', 'matter exchange', 'it cannot exchange matter', 'no matter could be transferred', 'addition of matter to the phase from without', 'flow of matter', and 'transfer of matter', in places where the poster of the edit would expect 'mass transfer'.
Page 7:
Page 9:
Page 13:
Page 16:
Page 18:
Page 51:
Page 108:
Page 117:
Page 119:
Page 121:
Page 129:
Page 145:
Page 155:
Page 203:
The enthalpy, H(S[P],P,{Ni}), expresses the thermodynamics of a system in the energy-language, or in the energy representation. As a function of state, its arguments include both one intensive (P) and several extensive (S[P],{Ni}) state variables. The state variables S[P], P, and {Ni} are said to be the natural state variables in this representation. They are suitable for describing processes in which they are experimentally controlled.
Initially, it will simplify the present account to rule that {Ni} be permanently fixed, so that one is considering a closed system, for which H = H(S[P],P) is the natural description of state. Such a system admits transfer of energy only as thermodynamic work and as heat, but not transfer of matter. This is the usual form of system considered in initial pedagogic accounts of thermodynamic work and heat.
For some problems, it is especially convenient to use enthalpy as the cardinal energy variable of the system.
This is so when the processes for the system are required to occur due only to thermodynamic operations of change of pressure in the surroundings, without heat transfer and without matter transfer. Matter transfer is prevented by enclosing the system with walls that are impermeable to matter, or by defining the moving boundary of the system so that no matter transfer occurs; such a system is said to be closed; its processes are transfers only of energy, as work and as heat. To consider a conceptual and idealized process in which there is negligible energy transfer into the system through friction within the system, so that heat transfer is entirely prevented, the processes of interest are conceptually further restricted to being infinitely slow (thus 'quasi-static', also called 'quasi-equilibrium'). Then the system can change volume only through thermodynamic work. The pressure, volume, and temperature of the system can then change. Changes of system volume and temperature are determined by the pressure changes.
It is then convenient to describe the system with enthalpy as the cardinal energy variable. Then one writes
In this way, the values of S[P] and P define the state of the system.
When it is convenient to consider that heat transfer shall occur at constant pressure P has been chosen as an independent variable becuais made when it is convenient to consider that thermodynamic work shall occur at constant (time-invariant) pressure, set by devices in the surroundings of the system. Then then thermodynamic work can occur only with changes of "V", the system volume. The system does thermodynamic work W = P ΔV on its surroundings when, in a thermodynamic operation, the pressure applied by the surroundings is changed by ΔP. In the sign convention used in physics (the contrary convention is used in chemistry) when ΔP < 0, then W > 0, and when ΔP > 0, then W < 0. In these circumstances, in a turn of language, it may be said that the piston is "permeable" to volume transfer under controlled pressure (transfer of V, the relevant extensive variable, with P as its conjugate intensive variable). The thermodynamic operation has imposed a change in pressure, the driver ΔP, causing a response ΔV, the volume that has been "transferred". End of nonsense.} This is the basis of the so-called adiabatic approximation that is used in meteorology.[1]
Alongside the enthalpy, with these arguments, the other cardinal function of state of a thermodynamic system is its entropy, as a function, S[p](H,p,{Ni}), of the same list of variables of state, except that the entropy, S[p], is replaced in the list by the enthalpy, H. It expresses the entropy representation. The state variables H, p, and {Ni} are said to be the natural state variables in this representation. They are suitable for describing processes in which they are experimentally controlled. For example, H and p can be controlled by allowing heat transfer, and by varying only the external pressure on the piston that sets the volume of the system.[2][3][4]
A temperature expresses hot and cold, as measured with a thermometer. There are various temperature scales that nearly or approximately agree with one another, but differ slightly because of the various characteristics of particular thermometric substances. The most commonly used scales are the Celsius scale (formerly called centigrade) (denoted °C), Fahrenheit scale (denoted °F), and Kelvin scale (denoted K). The kelvin (the word is spelled with a lower-case k) is the unit of temperature in the International System of Units (SI).
In physics, hotness is a body's ability to impart energy as heat to another body that is colder. In a body in which there are processes of chemical reaction and flow of matter, temperature may vary over its parts, and over time, as measured by a suitably small and rapidly responding thermometer, and may depend also on the match of the processes to the characteristics of the thermometer.
When a body has no macroscopic chemical reactions or flows of matter or energy, it is said to be in its own internal state of thermodynamic equilibrium. Its temperature is uniform in space and unchanging in time. Then, referring to the Boltzmann constant, a temperature scale is defined and said to be absolute because it is independent of the characteristics of particular thermometric substances and thermometer mechanisms. This is the Kelvin or thermodynamic scale, widely used in science and technology. The thermodynamic temperature is always positive, relative to an absolute zero.
For some conditions other than thermodynamic equilbrium, a suitable thermometer calibrated on the thermodynamic scale, including absolute zero, can register a negative temperature; such conditions are hotter than absolute zero.
Temperature is important in all fields of natural science, including physics, chemistry, Earth science, medicine, and biology, as well as most aspects of daily life.
A new revision of the lead has been posted here. The edit summary reads "The language here has been deteriorating; where is hotness defined like this? the temperature scales paragraph makes no accessible sense by its verbiage.)" No attempt was made to talk about this on the article's talk page.
I will content myself with a couple of observations. The first sentence of the revision, "Temperature is a physical quantity that expresses the subjective sensations of hot and cold", verges on nonsense. Subjective sensations are no so expressible. That is why they are called subjective. That is why temperature is defined as a physical quantity, not as a subjective sensation. There are plenty of textbook physical definitions of hotness, that have been cited in these pages, perhaps removed by other editors, apparently unnoticed by the reviser.
There are various statements of principle that are, by various writers, each respectively described as an expression of the second law of thermodynamics. They may be assembled under several headings. Several terms are useful for this purpose. A simple thermodynamic system is a body contained by walls of respective specified permeabilities, defined by its own internal state of thermodynamic equilibrium. A compound thermodynamic system is a collection of bodies separated by walls of respective specified permeabilities, defined by their mutual state of thermodynamic equilibrium. The surroundings of thermkdynamic system are not required to be thermodynamic systems, but may include one or more of them.
The statements by Carnot, Clausius, Kelvin, Planck.
Planck
Carathéodory
Planck
In theory, there is a unique absolute extreme of coldness, in which a body has only zero-point energy; according to the Third Law of Thermodynamics, its temperature, known as absolute zero, is approachable but unattainable through any actual physical process. It is denoted as 0 K on the Kelvin scale. In an ideal gas, and in other theoretically understood bodies, the thermodynamic temperature is proportional to the average kinetic energy of microscopic particles, which can be measured by suitable techniques. The proportionality constant is the Boltzmann constant; its value is defined by international convention; this defines the magnitude of the kelvin, the unit of the thermodynamic temperature scale.
Theoretically-based temperature scales are based directly on theoretical arguments, especially those of kinetic theory, thermodynamics, and quantum mechanics. They are more or less ideally realised in practically feasible physical devices and materials. Theoretically based temperature scales are used to provide calibrating standards for practical empirically based thermometers.
In physics, the internationally agreed conventional temperature scale is called the Kelvin scale. It is calibrated through the internationally agreed and prescribed value of the Boltzmann constant,[1][2] referring to motions of microscopic particles, such as atoms, molecules, and electrons, constituent in the body whose temperature is to be measured. In contrast with the thermodynamic temperature scale invented by Kelvin, the presently conventional Kelvin temperature is not defined through comparison with the temperature of a reference state of a standard body, nor in terms of macroscopic thermodynamics.
In theory, there is a unique absolute extreme of coldness, in which a body has only zero-point energy. According to the third law of thermodynamics, its temperature, known as absolute zero, is closely approachable but eventually unattainable through any actual physical process. It is used as a zero reference for the Kelvin scale, denoted as 0 K.
Apart from the absolute zero of temperature, Kelvin temperature of a body in a state of internal thermodynamic equilibrium is defined by measurements of suitably chosen of its physical properties, such as have precisely known theoretical explanations in terms of the Boltzmann constant. That constant refers to chosen kinds of motion of microscopic particles in the constitution of the body. In those kinds of motion, the particles move individually, without mutual interaction. Such motions are typically interrupted by inter-particle collisions, but for temperature measurement, the motions are chosen so that, between collisions, the non-interactive segments of their trajectories are known to be accessible to accurate measurement.
In an ideal gas, and in other theoretically understood bodies, the Kelvin temperature is defined to be proportional to the average kinetic energy of non-interactively moving microscopic particles, which can be measured by suitable techniques. The proportionality constant is a simple multiple of the Boltzmann constant. If molecules, atoms, or electrons,[3][4] are emitted from a material and their velocities are measured, the spectrum of their velocities often nearly obeys a theoretical law called the Maxwell–Boltzmann distribution, which gives a well-founded measurement of temperatures for which the law holds.[5] There have not yet been successful experiments of this same kind that directly use the Fermi–Dirac distribution for thermometry, but perhaps that will be achieved in future.[6]
The speed of sound in a gas can be calculated theoretically from the molecular character of the gas, from its temperature and pressure, and from the value of Boltzmann's constant. For a gas of known molecular character and pressure, this provides a relation between temperature and Boltzmann's constant. Those quantities can be known or measured more precisely than can the thermodynamic variables that define the state of a sample of water at its triple point. Consequently, taking the value of Boltzmann's constant as a primarily defined reference of exactly defined value, a measurement of the speed of sound can provide a more precise measurement of the temperature of the gas.[7]
Measurement of the spectrum of electromagnetic radiation from an ideal three-dimensional black body can provide an accurate temperature measurement because the frequency of maximum spectral radiance of black-body radiation is directly proportional to the temperature of the black body; this is known as Wien's displacement law and has a theoretical explanation in Planck's law and the Bose–Einstein law.
Measurement of the spectrum of noise-power produced by an electrical resistor can also provide an accurate temperature measurement. The resistor has two terminals and is in effect a one-dimensional body. The Bose-Einstein law for this case indicates that the noise-power is directly proportional to the temperature of the resistor and to the value of its resistance and to the noise band-width. In a given frequency band, the noise-power has equal contributions from every frequency and is called Johnson noise. If the value of the resistance is known then the temperature can be found.[8][9]
Historically, till May 2019, the definition of the Kelvin scale was that invented by Kelvin, based on a ratio of quantities of energy in processes in an ideal Carnot engine, entirely in terms of macroscopic thermodynamkcs. That Carnot engine was to work between two temperatures, that of the body whose temperature was to be measured, and a reference, that of a body at the temperature of the triple point of water. Then the reference temperature, that of the triple point, was defined to be exactly 273.16 K. Since May 2019, that value has not been fixed by definition, but is to be measured through microscopic phenomena, involving the Boltzmann constant, as described above. The new definition does not have a reference temperature.
An ideal material on which a macroscopically defined temperature scale might be based is the ideal gas. The pressure exerted by a fixed volume and mass of an ideal gas is directly proportional to its temperature. Some natural gases show so nearly ideal properties over suitable temperature ranges that they can be used for thermometry; this was important during the development of thermodynamics and is still of practical importance today.[10][11] The ideal gas thermometer is, however, not theoretically perfect for thermodynamics. This is because the entropy of an ideal gas at its absolute zero of temperature is not a positive semi-definite quantity, which puts the gas in violation of the third law of thermodynamics. The physical reason is that the ideal gas law, exactly read, refers to the limit of infinitely high temperature and zero pressure; these conditions guarantee non-interactive motions of the constituent molecules.[12][13][14]
In this article, the definition of quantity of energy transferred as heat is not the only definition used by reliable sources, but it is a distillation of what Wikipedia editors have found most consistent or widely agreed amongst an array of many different reliable sources. An attempt to state word-for-word all the definitions found in reliable sources would be like an attempt to state all possible views about the length of a piece of string, and is not undertaken in this article.
Working with the definition of quantity of energy transferred as heat in this article, one may observe that several kind of transfer are admitted. The primary definition in this article refers just to the energy received as heat by the thermodynamic system of interest, and does not specify how the energy leaves the surroundings of the system. The present definition admits two kinds of heat reception.
Some reliable sources speak of "conversion of work into heat" or "into heat energy" or "into thermal energy".[15][16][17][18] Such locutions refer to energy in transfer, not to energy in a static state of either system or surroundings by themselves considered separately. A few, especially older, reliable sources speak of "heat production" in these circumstances, but the term 'conversion' is more compatible with the present definition of transfer of energy as heat. Also, some reliable sources speak of "conversion of heat into work", or "thermal energy to mechanical energy";[19] the present definition of transfer of energy as heat does not use such locutions. This is because, even fictively, reversal of friction to provide work is not only physically utterly impossible, but is almost inconceivable. Planck regarded this fact as a prime example of the second law of thermodynamics.
These considerations may be expressed in mathematical terms for a closed system under controlled pressure, described in the energy language by H = H(S[P],P). One may observe that work done by factors or agencies in the surroundings of the system is not explicitly described in this formulation. That is to say, for example, that some pressure–volume work AP–V might be done on the system, and also that some isochoric work Aisoch might be done on it, defined within the surroundings. The present sign convention (that used in physics, opposite to that used in chemistry) counts thermodynamic work done by the system on its surroundings as positive, and thermodynamic work done by the surroundings on the system as negative. Isochoric work includes work that is simply measurable in the surroundings, such as for example shaft work. It also includes other kinds of frictional work, that are not easily, or even not at all, assessable in the surroundings. For example, when a body is compressed, its volume change incurs friction internally in its substance, depending on the speed, and perhaps the oscillatory character, of the compression. Such frictional work cannot be directly assessed in the surroundings. It has to be inferred, by use of the First Law, and from knowledge of the thermodynamic work done in the complete process. That relies on knowledge, acquired in extensive experiment, of the thermodynamic properties of the substance of the body or thermodynamic system of interest. For example, one may consider a process in which initial and final volumes differ by ΔV = Vfinal − Vinitial < 0, while the initial and final pressures are equal, P. Then the thermodynamic work done by the system is W = P ΔV < 0.
Then, as defined in the surroundings, the work done on the system is
so that
As defined for the thermodynamic system of interest, the thermodynamic work done by it is then
Then, in this language, in the present sign convention, as usual, the first law of thermodynamics will be written
From this, by substitution, one gets
If one wishes to write these formulas in terms of the work A, as defined in the surroundings, done by the surroundings on the system, and the quantity Q of energy transferred to the system by mechanisms other than thermodynamic work and transfer of matter, and in terms of the quantity Qcon-rad of energy transferred by conduction and radiation to the system from the surroundings, one will write
Then, continuing to write with the term A for the work done by the surroundings on the system, and eliminating ΔH, from (vii) and (viii), one has
Thence
so that, with the present definition of heat as quantity of energy transferred to the system by mechanisms other than thermodynamic work and transfer of matter, one has
Taking up the admirable comment that it is not easy to find reliable sources for our presentation of the four laws.
At the risk of being seen as excessively garrulous, I first note that we had a lot of rain here last night and that the river is running very high. I am commenting for information only, no action intended.
Moreover, I have had another look at the thoughts of Callen, who does not present the laws in precisely the terms 0, I, II, and III. He offers his own system of postulates, which are nearly or practically the same as the usual four laws.
I am wishing particularly to chat about his Postulate IV. He writes
In a much later chapter of his book, he writes
As I read this, it means that in a simple thermodynamic context, the Planck form relies on statistical mechanical notions, beyond classical thermodynamics per se.
Guggenheim is circumspect in stating the third law. He concludes
I will not pursue his view further here.
Adkins writes
He then proceeds to derive
For information only, no action intended.
I like Callen's account of the Legendre transforms of system energy and of system entropy. An example that comes to mind is enthalpy. The only source that I know that uses the term 'cardinal function of state' is Tschoegl. He uses it to qualify internal energy and its corresponding entropy. One might say that the one-source character makes the term idiosyncratic. I like the term, to signify that the internal energy and its corresponding entropy describe an isolated state if the extensive variables are held fixed, with no transfers permitted. A transfer is specified by a change in an extensive variable.
I like to think of enthalpy as a function of its natural state variables, such as {entropy, pressure} for a closed system. Then it would have a reference state defined by a distinguished {entropy, pressure} value. A problem is to relate such a quantity to internal energy, which has a reference state defined by a distinguished {entropy, volume} value. To define enthalpy, IUPAC proposes the Euler relation . Presumably that means that one defines a common reference state as for internal energy, with a reference {entropy, volume} value, and measures the pressure at that reference state, and uses that pressure to define the reference state for enthalpy.
Starting at that reference state and using adiabatic walls to prevent heat transfer, if then, slowly enough to annul the kinetic energy imparted by the system to the surroundings, one reduces the pressure of the surroundings, to get a new state, then the system will expand, doing work on the surroundings, and reaching a new volume , where , and a new pressure, , that of the new surroundings, where . I think this guarantees an adiabatic process. The entropy will not change. The energy will increase by .
In a second process, one may then make a system wall diathermic, and slowly enough heat the system by conduction, till its temperature reaches that of the reference state.
Let us define
These can be solved to give us
Let us stipulate a common reference state in which the body has temperature , pressure , and volume , which we can measure without knowing the energy or entropy.
We are free to define for the common reference state the values and .
We are free to stipulate
that , with , and
that , with .
We are still free to stipulate the common reference state values .
Consequently, we are stipulating that
We also have
It can be convenient for purposes of exposition to use the term 'intrinsic energy of the system'. The term covers state functions such as internal energy, enthalpy, Helmholtz free energy, Gibbs free energy, as well as others. Here it is convenient to use the symbol to denote an intrinsic energy. As a function of state, that symbol is incomplete, lacking definition of the dummy argument (.).
This notation may be exemplified by writing
where now is the customary notation for the customarily named state function 'internal energy', with all its arguments being extensive state variables.
Another example is
where now the intrinsic energy is the customarily named state function 'enthalpy', with its arguments as shown, just one of them, , being an intensive state variable.
Another example is
where now the intrinsic energy is the customarily named state function 'Helmholtz free energy', with its arguments as shown, just one of them, , being an intensive state variable.
A further example is
where now the intrinsic energy is the customarily named state function 'Gibbs free energy', with its arguments as shown, just two of them, and , being intensive state variables.
At this stage of this account, the list of customary names is running short, but in principle there are more possible versions of intrinsic energy. In order for the intrinsic energy to be a characteristic function of state, at least one of the argument state variables must be extensive.
An example of an intrinsic energy, lacking a conventional name, is
where now just one of its arguments, , is an intensive state variable.
Callen offers another convenient systematic notation for the present topic. It distinguishes the intensive state variable arguments, leaving unmentioned the remaining extensive state variables.
For example, where the foregoing notation writes
Callen writes
Enthalpy is a mathematical abstraction that keeps account of transfers of energy that befall a body of radiation and matter that has a chemical composition defined by a suitable list of chemical constituents. To understand this, some background knowledge of thermodynamics is necessary. The present account covers only a simplified range of thermodynamic variables and processes, for a simple system.
With two exceptions that epitomise thermodynamics, variables and functions of thermodynamic state are directly determined by respective single concrete measurements, for example temperature and mass. The thermodynamic state of a body is always fully specified by the values of several pairs of such measurements, along with temperature. For example, the volume and pressure, and the mole numbers and chemical potentials of the chemical constituents, along with the temperature. It is remarkable and notable, and a most basic fact of nature, and a foundational postulate of thermodynamics, that bodies in such states of matter and radiation exist, and persist or endure over time if the body is isolated. It is sometimes called the minus-one-th law of thermodynamics. They are called states of thermodynamic equilibrium.
Taking into account full knowledge of the specific properties of the chemical constituents, and observing that it refers to a state of thermodynamic equilibrium, such a specification uses one more single-measurement quantity than is necessary. A prime example is that temperature can be omitted from the list, while taking into account the specific properties of the chemical constituents.
Thermodynamics works by using mathematical abstractions that summarise the specific properties of the chemical constituents. In general terms, these abstractions are (intrinsic) energy and entropy. These are the two exceptions mentioned just above. They can be specified in what are known as 'characteristic functions', and in what are known as 'fundamental equations', to be formulated below. As summaries of much detailed information, they can be determined only by a range of several or many concrete measurements, and are defined with respect to particular kinds of thermodynamic process. They are called 'characteristic' because they summarise the particular physical and chemical properties that uniquely characterise the listed chemical constituents, and 'fundamental' because they provide a full summary of all relevant thermodynamic information. The fundamental equations or characteristic functions provide information that is not fully present in any single equation of state, but they can be used to construct several various equations of state.
Of the characteristic function pairs, two are distinguished, and may be called the cardinal functions. They are functions only of extensive variables of state, not including intensive variables of state. They are called internal energy and entropy; for clarity and brevity in this section of this article, the latter will be called cardinal entropy. Enthalpy is a characteristic function that is convenient for describing certain kinds of thermodynamic process that are not conveniently described by internal energy and cardinal entropy.
To understand enthalpy, it is convenient to think initially about the cardinal pair of what are known as the Euler relations. They presuppose knowledge of internal energy and cardinal entropy, and are formulated with respect to a chosen reference state of the body of interest, that fixes certain reference constants:
Beyond observing that the Euler relations express much about the fundamental concepts of thermodynamics, explaining their logic and source is beyond the scope of the present account.
As functions of state, internal energy and cardinal entropy refer to processes defined by transfers of prescribed amounts (including zero amounts) of extensive variables, such as cardinal entropy, volume, and mole number.
For describing processes defined in the same way except that, instead of an amount of volume transferred, a pressure imposed from the surrounding is prescribed, a particularly convenient formulation is provided by enthalpy, . The corresponding formulas are
Ideally, the processes that define them are, for internal energy, transfers of energy, as thermodynamic work, and of matter, and for cardinal entropy, transfers, of energy, as heat, and of matter. In terms of mathematical symbols, this may be written:
where denotes internal energy, denotes cardinal entropy, denotes system volume, and denotes the set of mole numbers that constitute the system. The pairing of internal energy with cardinal entropy indicates that each of them is a monotonic function of its arguments, and that each can be obtained by solving the other.
For closed systems, in which matter transfers are prevented, the thermodynamic processes of interest may be formulated:
where denotes the increment in internal energy due to the system's having done thermodynamic work, <>W</math>, by a prescribed expansion against a resisting pressure in the surroundings, and having gained energy as heat from the surroundings, and denotes the increment in cardinal entropy due to those transfers.
Instead of choosing that work done by the system be determined by a prescribed change of volume, it may be convenient to let the work done be determined by holding the surrounding resisting pressure at a prescribed value. For such processes, another special pair of characteristic functions are necessary. Instead of 'internal energy', the intrinsic energy is then called enthalpy. There is no customary special name for the corresponding entropy, but for clarity in this section of this article, it will be called enthalpic entropy.
The state variables and functions are then reformulated:
where denotes enthalpy, denotes enthalpic entropy, denotes system pressure, and denotes the set of mole numbers that constitute the system.
For closed systems, in which matter transfers are prevented, the thermodynamic processes of interest may be formulated:
where denotes the increment in enthalpy due to the system's having done thermodynamic work, by some measured expansion against the prescribed resisting pressure in the surroundings, and having gained energy as heat from the surroundings, and denotes the increment in enthalpic entropy due to that gain of enthalpy.
Boltzmann defined two kinds of monode, the ergodes and the holodes.
The 1980 article by Smith[1] has occurred to my drifting mind.
I have become slightly less ignorant since I last considered that article. I now have, I think, a clearer view, in the light of Born's comments on the topic.
My current view is that heat and work can be uniquely distinguished primarily for processes in closed systems (matter transfer not permitted), but for open systems only when the heat transfer path and the work transfer mechanism/path are between the system and its surroundings and are distinct from every matter transfer path.
It is conceptually key that the system and surroundings, and transfers between them, be uneqivocally distinguished and defined.
For uniqueness, it is practically necessary to restrict attention to spontaneous or natural processes. In practice, this means letting the system do thermodynamic work on its surroundings while keeping the heat and matter transfers separate from the work transfer. Then the heat and matter energy transfers can also be uniquely defined. The thermodynamic work transfer can in principle be measured as gain in internal energy of an otherwise isolated compartment in the surroundings, provided that the work be done monotonically. For example, the system can be allowed to expand against pressure exerted on it from the surroundings through a piston, during the process the surroundings pressure being kept only a little less than the system pressure. During the work transfer, heat and matter transfers are also permitted, through their separate pathways. If there is matter transfer, then the energy transferred in it must be uniquely defined, between the system and a compartment of the surroundings that is otherwise isolated; in other words, that compartment is permitted to transfer neither energy nor matter to or from any other part of the surroundings.
[https://en.wikipedia.org/w/index.php?title=Entropy&diff=next&oldid=84833493 This edit, dated 31 October 2006, and still standing, added the following paragraph.
{{cite book}}
: Italic or bold markup not allowed in: |publisher=
(help)Sad to say, I can't work out what first sentence of this means.Chjoaygame (talk) 10:53, 11 March 2021 (UTC)
The characteristic extensive thermodynamic functions of state are energy and entropy, specified in various particular ways. For convenience of exposition, considered first here are the various energy functions, which are Legendre transforms of one another, known as thermodynamic potentials, several of which have well known names, such as 'the Gibbs free energy'. Further along below, the various entropy functions are considered; they are known as Massieu functions.
The cardinal energy state function is the internal energy, specified as a function of the entropy and of suitable ordinary physical variables, such as volume. For example, a Legendre transformation converts the internal energy into another energy function or thermodynamic potential (called enthalpy) of entropy and pressure, where the pressure is the conjugate state variable of the volume. A different Legendre transform converts the internal energy into another energy function, called the Helmholtz free energy, switching the argument entropy to the argument temperature, which is the state variable conjugate to the entropy. One of a conjugate pair is extensive, the other intensive.
Newton, in 1721, in his Opticks, in a long discussion of bodies and motions, elastic and inelastic, on pages 350 and 373, wrote:
We are looking at writings in several languages, from several historical epochs. A glossary may help.
It is no simple thing to say how vis viva is to be seen as 'belonging' to a 'body'. It needs to be said how a 'body' is constituted, and how it can have intrinsic and relational properties. It was not initially clear whether to regard such a quantity as as intrinsic or relational. Leibniz regarded space as relational.
Over a collision between two bodies and , if the collision is elastic, then the sum is unchanged. If the collision is inelastic, it is not. But perhaps the missing quantity of vis viva has passed into microscopic components, Newton's "small Particles" of the bodies. This needs to be precisely articulated, and was controversial. According to Alejandro Jenkins "Leibniz argued that in an inelastic collision the missing vis viva is transferred to the imperceptible motion of the bodies’ microscopic components. This is consistent with our current understanding of dissipation, but at the time it was an untestable speculation." Loemker writes of Leibniz's expression of "the principle of the conservation of force (motive action)—that in any system of moving bodies the sum of is a constant." Also Leibniz knew of conservation of 'quantity of progress', our 'momentum'. Loemker seems to say that Leibniz used the terms conatus, momentum, and vis viva, for , , and respectively, but examination of the text seems to refute this, at least in part. It seems that Loemker's thinking may be muddled?
Newton knew that over a right line collision between bodies, elastic or not, the sum is unchanged. Momentum is conserved even in an inelastic collision.
The following, with direct quotes, is from Alejandro Jenkins (https://arxiv.org/abs/1301.3097v2).
In 1644, Descartes proposed that a positive scalar quantity is conserved, where the are the positive speeds of the bodies. This was proposed as a law of conservation of 'quantity of motion'. Obviously, it is a law of conservation neither of vector momentum, nor of scalar energy.
"In papers presented before the Royal Society in 1668, John Wallis and Christopher Wren showed that the quantity conserved upon a collision was" the vectorial sum .
"Shortly afterwards, Christiaan Huygens published a result that he had already derived in the 1650s: that elastic collisions conserve another quantity as well, corresponding to twice the modern kinetic energy, ."
Mittelstrass & Aiton:
Bertrand Russell:
Daniel Garber (lots here):
Du Châtelet’s book Foundations of Physics was published in 1740. Besides her analysis of the physics of heat, light, and fire, and her translation of Newton, it expresses her main contributions to rational mechanics.
In §579, du Châtelet states that Leibniz was the inventor of the concept of "forces vives" [in Leibniz's Latin vires vivae]. The vis viva of a body of mass moving with a speed is measured by . This is distinct from that body's "force morte", measured by , also called by Leibniz "quantity of motion".
Leibniz had discovered the law of conservation of momentum. Du Châtelet observes that Leibniz was the first to distinguish between "force vive" and "force morte".) Leibniz is translated (Ariew & Garber p. 157) as writing in "On nature itself" (1698):
Leibniz (A & G p. 50) proposes a distinction between "force" and "quantity of motion" as follows. "Force" in this sense is measured as directly proportional to , where denotes the mass of a body lifted through a height . Assuming a conservation law, Leibniz lets the bodies of masses , fall back through heights and , reaching speeds and . They attain quantities of motion and . Leibniz proposes that though the "forces" and be made equal, the "quantities of motion" and will in general be unequal, thus distinguishing "force" from "quantity of motion". At this point in this letter, he does not mention the quantities and .
Elsewhere, (Loemker p. 299) Leibniz writes "the living forces of equal bodies are not proportional to their velocities but to the squares oftheir velocities." Also (A & G p. 111): "the powers of unequal bodies are jointly proportional to the size of the bodies and the square of the velocities."
According to Ducheyne, the Leiden physicist 's Gravesande was a respected expositor of Newtonian physics (natural philosophy). For example, Voltaire went to Leiden to discuss it with him.[1]
Quote from Brush: "The term potential energy was introduced by Rankine around this time and immediately adopted by Thomson. W. J. M. Rankine, Proc. Glasgow Phil. Soc. 3, 276 (1853); Papers, p. 203; footnote on p. 554 of Thomson's Papers 1."[2]
Editor ReyHahn, thank you for your care in citing texts. I much respect that.
Editor Kbrose, thank you for your comment.
I will try to reply with two wings. One will be a more direct reply to the comments of Editor ReyHahn, the other to those of both Editors just mentioned, in a more general approach.
Looking at the sources listed by Editor ReyHahn:
Some of them are not standard texts on thermodynamics.
Deffner & Campbell. I regret that still don't have access to this. I would point out that it is not primarily about thermodynamics as such.
Nag. I have access to the 2002 reprinting of the first edition and to the second edition (2010), not the 2008 14th reprinting of the first edition that you link to. For the second edition, the author removed the chapters that cover the Onsager relations. One might say that this counts against the Onsager relations being regarded as belonging to thermodynamics as such.
Tiwari has a sense of humour that I like. On page 542 he writes:
I don't accept the idea that Tiwari attributes to Nernst, that Rankine should not appear in a short list of originators of the second law.
I think Tiwari's book isn't primarily about thermodynamics as such, and doesn't justify suggesting in our article that the Onsager relations might be a fourth law of thermodynamics.
Srivastava & Saha & Jain write on page 214:
These authors feel that so-called "non-equilibrium thermodynamics" makes classical thermodynamics obsolete. With respect, judging from your having cited them without critical comment, I guess that you might agree with them. Please tell me that my guess is mistaken.
My opinion is that the paragraph from Srivastava et al. that I have copied just above is riven with nonsense. The authors are misled in precisely the way that gives me reason to oppose putting in material that suggests that the Onsager relations might be laws of thermodynamics. I guess that you might think I am misled in this.
The nearest that I can offer in reply is to refer to Phil Attard's concept of many-time entropies. I wrote to him, asking how to measure them. He replied that he didn't know. To me, this implies that they do not belong to thermodynamics in the ordinary sense of the word. In a nutshell, Srivastava et al. is not a Wikipedia reliable source on this point.
I am not persuaded by the Nobel Prize speech.
Now looking more generally.
My reason is respect for clarity of thought. Thermodynamics is a monumentally clear and straightforward theory, for which Einstein gave a celebrated distinction. It is essentially macroscopic, and builds upon the distinction between heat and work, to construct its key quantity, entropy. So-called 'non-equilibrium thermodynamics' is a radically different topic, resting on a direct contradiction of the foundations of thermodynamics. Wikipedia should not erase that. To erase it would be to mislead newcomer readers, likely incurably.
The heart of it is that entropy is well defined only for a thermodynamic system in its state of internal thermodynamic equilibrium. That is why it makes sense of Boltzmann's formula, which expresses a symmetry, with all 'states' equally weighted. This sets up the possibility of a logically sound statistical mechanical explanation of the second law.
It is regrettable to read a suggestion that 'non-equilibrium thermodynamics' is a more advanced topic. Yes, it is a more difficult topic. That is because it deals with more difficult problems. But it hasn't yet established an adequate generalization of the concept of entropy. That would make it truly advanced.
I have repeatedly said that the Onsager relations are important. But my main point is that they do not belong to thermodynamics as such. My above list of standard texts, that discuss them but do not say that they are called the fourth law of thermodynamics, is not a sample of thermodynamics texts in general. It is a sample of a smaller set, those that discuss the relations. Many standard texts on thermodynamics do not mention them, for example, Bailyn, M. (1994), A Survey of Thermodynamics, American Institute of Physics, New York, ISBN 0-88318-797-3. I think this counts against the Onsager relations being regarded as belonging to thermodynamics as such.
I think that 'non-equilibrium thermodynamics' is a misnomer, or clang word-association. It is about transport theory; it does not belong to thermodynamics. Yes, it uses the term 'entropy production', but it doesn't actually use the concept to calculate entropy because it can't. Alternatively, one might regard 'non-equilibrium thermodynamics' as an approximation, though without clearly specified scope, not thermodynamics as such. The calculation of time rate of entropy production is interesting, but it is generally observed that it does not contribute to transport calculations, which are based on transport coefficients that do not use the concept entropy. These facts are not emphasised by many texts on so-called 'non-equilibrium thermodynamics'. When it came out, I was thrilled to read Glansdorff & Prigogine's Thermodynamic Theory of Structure, Stability, and Fluctuations (1971), Wiley–Interscience, London. But that did not make it belong to thermodynamics as such; though it discusses them, it does not suggest that the Onsager relations should be called a fourth law of thermodynamics, so that it could be added to my list. Though I recognise that I may not persuade you about it, I think that putting the Onsager relations in this article would detract from the quality of Wikipedia.
Page 230:
Page 400:
Page 151:
Page P-1
Page 2:
Page 3:
Page 302:
The article on heat was wrecked on the impetus of a quote from Steven Weinberg's 2021 book Foundations of Modern Physics, which was treated as a reliable source. Weinberg may be an expert on modern physics, especially on quantum physics, but he is careless and slipshod on the topic of heat in thermodynamics, disqualifying himself as a reliable source on it. Weinberg is inconsistent in slipping between 'heat' as in ordinary language and as a technical term in thermodynamics. This is partly due to Weinberg's slipping between statistical mechanics and thermodynamics, which will tend to confuse a reader who is new to the topics.
Weinberg's book's section 2.1, 'Heat and Energy', derives a law of conservation of kinetic energy in a wrong way. It examines a collision on the basis that the colliding bodies interact only when they are in contact, forgetting that the proper basis is that the collision should be elastic. An inelastic collision converts kinetic energy into heat imparted to the inelastic body, not by virtue of a temperature difference.
Section 2.2, 'Absolute Temperature', carelessly says that the first law of thermodynamics is the conservation of energy, without mentioning heat and work. The section goes on to say that the 1850 paper by Clausius shows that it is possible to find a definition of temperature with absolute significance by the study of thermodynamic engines known as Carnot cycles. The section then goes on (without mention of Kelvin) to develop Kelvin's thermodynamic definition of absolute temperature, taking the development from Fermi's little 1936 book Thermodynamics. Reading the 1850 paper of Clausius, contrary to Weinberg's assertion, I think one can hardly find thermodynamic temperature defined by a ratio of quantities of heat in a Carnot cycle, as in Fermi's book and as proposed by Kelvin.
In ordinary language,[1] heat is what makes a thing hot.
The rest of this article concerns thermodynamics,[2] for which heat is energy in transfer to or from a thermodynamic system, by mechanisms other than thermodynamic work or transfer of matter.[3][4][5][6][7][8][9] The various mechanisms of energy transfer that define heat are stated in the next section of this article.
Like thermodynamic work, heat transfer is a process involving more than one system, not belonging to any one system. In thermodynamics, energy transferred as heat contributes to change in the system's cardinal energy variable of state, its internal energy. This is to be distinguished from the ordinary language conception of heat as belonging to single system.
The measurement of energy transferred as heat, performed by measuring its effect on the states of interacting bodies, is called calorimetry. For example, heat can be measured by the amount of ice melted, or by change in temperature of a body in the surroundings of the system.[10]
This edit https://en.wikipedia.org/w/index.php?title=Heat&diff=next&oldid=1114283454 was made at 2022 on 8 Oct 2022 by anonymous IP editor 99.113.71.3 .
The mechanisms of energy transfer that define heat include conduction, through direct contact of immobile bodies, or through a wall or barrier that is impermeable to matter; or radiation between separated bodies; or friction due to isochoric mechanical or electrical or magnetic or gravitational work done by the surroundings on the system of interest, such as Joule heating due to an electric current driven through the system of interest by an external system, or through a magnetic stirrer. When there is a suitable path between two systems with different temperatures, heat transfer occurs necessarily, immediately, and spontaneously from the hotter to the colder system. Thermal conduction occurs by the stochastic (random) motion of microscopic particles (such as atoms or molecules). In contrast, thermodynamic work is defined by mechanisms that act macroscopically and directly on the system's whole-body state variables; for example, change of the system's volume through a piston's motion with externally measurable force; or change of the system's internal electric polarization through an externally measurable change in electric field. The definition of heat transfer does not require that the process be in any sense smooth. For example, a bolt of lightning may transfer heat to a body.
Convective circulation allows one body to heat another, through an intermediate circulating fluid that carries energy from a boundary of one to a boundary of the other; the actual heat transfer is by conduction and radiation between the fluid and the respective bodies.[1][2][3] Convective circulation, though spontaneous, does not necessarily and immediately occur simply because of some slight temperature difference; for it to occur in a given arrangement of systems, there is a threshold that must be crossed.
Although heat flows spontaneously from a hotter body to a cooler one, it is possible to construct a heat pump which expends work to transfer energy from a colder body to a hotter body. In contrast, a heat engine reduces an existing temperature difference to supply work to another system. Another thermodynamic type of heat transfer device is an active heat spreader, which expends work to speed up transfer of energy to colder surroundings from a hotter body, for example a computer component.[4]
The mechanisms of energy transfer that define heat include
Convection allows one body to heat another, through an intermediate circulating fluid that carries energy from a boundary of one to a boundary of the other; the actual heat transfer is by conduction and radiation between the fluid and the respective bodies.[6]
When there is a suitable path between two bodies with different temperatures, heat transfer by conduction and radiation occurs necessarily, immediately, and spontaneously from the hotter to the colder system. Convection, though spontaneous, does not necessarily and immediately occur simply because of some slight temperature difference; for it to occur in a given arrangement of systems, there is a threshold that must be crossed.[7]
Although heat flows spontaneously from a hotter body to a cooler one, it is possible to devise a heat pump which expends work from a source in the surroundings to transfer energy from a colder body to a hotter body. In contrast, a heat engine reduces an existing temperature difference to supply work to another system. Another thermodynamic type of heat transfer device is an active heat spreader, which expends work to speed up transfer of energy to colder surroundings from a hotter body, for example a computer component.[8]
Since the days of Heraclitus, who said that no man crosses the same river twice, and of Aristotle, who thought in terms of 'substances', philosophy has practically distinguished two kinds of ultimate ontological actual or natural entities, 'process' and 'substance'. It might be said that substances are subject to change, but processes are not, for they are change. Or that substances are changed by processes, and that processes change substances.
This is relevant to the word 'heat'. In ordinary language, one can speak of heat as a substance, while, in thermodynamics, where one can think of a thermodynamic system as a substance, heating is properly thought of as a process, so that the word 'heat' properly refers to a process as contrasted with a substance. This is the import of the overthrow of the caloric theory of heat, which rests entirely on calorimetry and a sort of 'conservation of heat'. In ordinary language and in thermodynamics, 'friction' refers to a process that produces heat.Chjoaygame (talk) 10:25, 2 July 2023 (UTC)
Calorimetry relies on conservation and dispersal of heat, while the mechanical theory of heat thinks of production of heat through dissipation of work or of potential energy, and it thinks of consumption of heat through spontaneous work. Chemical reactions convert chemical potential energy into internal energy. Friction converts mechanical energy into internal energy. Internal energy can supply work as well as heat.
A change of the internal energy of a body consists of three components:
Those components of change are not in general separately conserved. They jointly draw on or contribute to the common pool of internal energy. In a thermodynamic process, those components change in various proportions, so that
Here the symbol refers to the change in chemical potential energy, which is negative in a chemical reaction, signifying dissipation of chemical potential energy, but positive when matter enters the working body.
For fictive idealised processes, one can imagine the components to occur separately.
For a compound system consisting of two bodies and that are initially in thermodynamic equilibrium across an arrangement of walls, then subject to a thermodynamic operation that increases some permeabilities of the walls, in such a way that the bodies actually further interact, after the further interaction it always happens that
where and
This is one statement of the second law of thermodynamics.
Friction, in rubbing, in viscosity, in electrical conduction, and in hammering, is simply irreversible; one cannot undo friction. It was the turning point in physicists' escape from the caloric theory.
In friction, there is conversion of work into heat. The process is a transformation of energy as well as a transfer of energy. The arriving form of energy, in the body or thermodynamic system, is different from the departing form of energy, in the surroundings.
Some of the heat generated by friction can be recovered as work, but not by simple reversal of friction.
Adkins starts with a definition of temperature, partly based on the zeroth law. He uses the ideal gas scale, calling it 'thermodynamic temperature'. I wonder how Kelvin would feel about that? Buchdahl would call it an 'empirical temperature'.
Now, as to heat in Adkins, I didn't find what I could recognize as an explicitly 'introductory' definition. I found a general exposition of the move from the Laplace–Lavoisier doctrine of caloric to the current thermodynamical theory of heat.
One of the factors in the move was the cannon boring experiment of https://en.wikipedia.org/wiki/Benjamin_Thompson, in his article https://en.wikipedia.org/wiki/An_Experimental_Enquiry_Concerning_the_Source_of_the_Heat_which_is_Excited_by_Friction.
Thompson's idea was accepted by https://en.wikipedia.org/wiki/Julius_von_Mayer in 1842, as follows:
Mayer gave a calculation of the mechanical equivalent of heat, relying on frictional generation of heat in paper pulp, and on calorimetry.
A further factor in the move to the current theory was the water paddle experiment of https://en.wikipedia.org/wiki/James_Prescott_Joule, another version of heat generation by friction. That form of heat generation was examined by Joule, and described in his 'first law', https://en.wikipedia.org/wiki/Joule_heating.
Explicitly defining heat, Adkins writes on page 32:
No mention of temperature here. Apparently, for a closed system, heat is defined by exclusion of work, along with the postulate of conservation of energy.
Eventually, Adkins is not quite clear about his definition of heat. He writes:
Moreover, Adkins is not quite clear about his definition of work. I think his unease in decision is due to his apparent failure to be clear, for a closed system, about the difference between work defined solely in the surroundings, and thermodynamic work defined by changes in the thermodynamic system's ordinary physical state variables other than temperature or entropy. An energy transfer as thermodynamic work done by the system on its surroundings is defined by both the changes in its ordinary physical state variables other than temperature and entropy, and the ordinary physically defined force × distance work. Apart from the idealization of an infinitely slow process, an energy transfer as work done by the surroundings on the system, defined through ordinary physical work in the surroundings, involves friction in the system, and is not identical with the associated thermodynamic work. Adkins does not tell us how the work done by forces in the surroundings reaches the system without the operation of the second law, that observes that friction is likely to occur.
At this stage, Adkins is not considering transfer of matter.
Enough on Adkins for the moment.Chjoaygame (talk) 01:10, 4 June 2023 (UTC)
Anderson, G., 2005, Thermodynamics of Natural Systems, Cambridge University Press, is rather chatty in his introductory chapters. He writes:
Apparently, Anderson, despite all his care, has forgotten about friction as a source of heat. Also, he is rather vague about what he means by "in response to a force moving through a distance." What is the cause of the force? It makes a difference whether the force is generated by the system in a process of spontaneous expansion, or by some factor in the surroundings. The force itself passes the work energy, which is therefore not a response to the force but is an aspect of the force itself.
In his next chapter, about the first law of thermodynamics, Anderson goes on to talk about water entering and leaving a pond, in an analogy that is almost shamelessly chatty, and lacking in thermodynamic insight. He goes on to say:
Anderson goes on to talk about temperature without having given a thermodynamic definition of it. What he says is pretty much covered by what Buchdahl calls 'empirical temperature', and makes sense in that light. Anderson is largely concerned with chemical reactions and enthalpy. He is aiming to define "chemical energy". His actual words are in the first paragraph of his chapter on the first law:
That is perhaps all that Anderson thinks the student needs to know about thermodynamic temperature at that point. Anderson goes on to postulate a state variable called 'entropy'. He says of it:
Soon, Anderson says:
But he doesn't pursue this, saying that it is scarcely necessary for the purpose of his book.
Later, Anderson remarks:
Evidently, an electrical heater can be used to generate heat, as noted by Joule.Chjoaygame (talk) 09:15, 4 June 2023 (UTC)
Now to look at Principles of Thermodynamics by J.-P. Ansermet and S.D. Brechet, 2019, Cambridge University Press. They start with a picture of Joule, with a caption that says "In 1840, he stated the law that bears his name on power dissipated by a current passing through a resistance." Their actual text begins with 1839 work by Marc Séguin on the heat engine. They go on to mention that in 1842 Julius Robert von Mayer, in a treatise, asked "what is the change in temperature of a stone when it hits the ground after falling from a given height?" They remark that Joule actually measured that in 1845. I would say that the rise in temperature was due to friction of impact within the stone.
On the first law of thermodynamics, they say
Perhaps following Prigogine and Defay, they are happy to talk of rates of energy flow, which does not respect the rule of classical thermodynamics, which refers situations of changes of state from one thermodynamic equilibrium to another. I do not like the confusion that is introduced by talking of heat transfer as well as of heat "exchange". I think that talk of heat "exchange" is slippery. Max Born is careful to say that the energy transfer that accompanies matter transfer in itself cannot be resolved into heat transfer and work. The two latter forms of transfer, if simultaneous with matter transfer, to be identified, must occur by pathways separate from matter transfer. For example, radiative transfer of heat can be distinguished from energy transfer accompanying matter transfer. Maxwell said that matter transfer by convection was not a form of heat transfer. There is no mention of radiative transfer nor of friction here.
Ansermet & Brechet go on to remark that
Soon they remark that
In such a collision, the colliding bodies exhibit conversion of the kinetic energy of their relative motion into internal energy. Since there is a temperature increase in the bodies, one might be tempted to say that this occurred due to the internal friction in the collision. One might even consider talking of generation of heat. But Ansermet & Brechet talk only of kinetic energy and of internal energy, not of friction, nor of heat in this process. To say that the product is internal energy is to avoid saying whether it is transferred as heat or as work; that is equivocation like that of Adkins.
Ansermet & Brechet consider a fan in a room. They consider a viscous frictional torque associated with the motion of the fan, and calculate it by considering entropy production, but do not mention heat in this scenario. Again, to say that the product is entropy is to avoid saying whether the energy transfer is as heat or as work; that is equivocation like that of Adkins.
They consider a damped harmonic oscillator, and ask the student to calculate the power P(t) due to the friction force, but they do not use the word 'heat' in this scenario.
In a section on heat in their chapter on the second law, they write:
Maxwell and Born would have qualms about that. Benjamin Thompson and Planck would be unhappy that it does not mention friction or rubbing.
Ansermet & Brechet deal with what Benjamin Thompson and Planck would call generation of heat sometimes by talking of entropy production, other times of heat production. They write:
Enough about Ansermet & Brechet for the moment.Chjoaygame (talk) 10:56, 4 June 2023 (UTC)
Now to look at Atkins, P., de Paula, J., Keeler, J., 2018, Physical Chemistry, eleventh edition, Oxford University Press.
Their first sentence, for closed systems, that could be considered as a start to defining heat reads:
This expresses the admirable ideas of Bryan (1907) and of Carathéodory (1909) that are prime ingredients of rigorous modern thermodynamics.
In my opinion, Atkins et al. then go overboard by writing:
I agree that thermodynamic quantities are always measureable through the surroundings. But that includes the internal state variables of the system. Atkins et al. demand that all features of a process of transfer of energy are defined by considerations that exclude the externally measured internal state variables of the system. They dismiss friction and rubbing, which Planck considered important; sad to say, it isn't easy to find English translations of Planck saying so. Atkins et al. use the notion of 'organized motion of atoms in the surroundings'. That is a notion foreign to thermodynamics proper, and hard to define in simple terms. Pressure in the surroundings is easily thought of as a manifestation of disorganized motion of atoms or molecules, but is able to do work as defined in the surroundings. In a real process, however, it will cause friction within the system, so that not all that work as defined in the surroundings will reach the system as thermodynamic work, which, for its definition, requires consideration of the externally measured internal state variables of the system. This is an example of what Adkins means by loss of clarity of thought when he talks about thermodynamics as a strictly macroscopic theory, as in his above quote.
Atkins et al. go on to remark that
When a gas expands into a vacuum, there is transfer of matter from the originally enclosed body of gas to the originally empty space. The process is set going by the thermodynamic operation of removal of the partition between the two spaces. In the view of Max Born, such a process does not allow a distinction between heat and work because it includes transfer of matter. This is one reason why the first law is mainly about closed systems. This seems to be forgotten by Atkins et al..
It is becoming evident that I am inclined to prefer the thinking of the best and most reliable sources, as distinct from less reliable sources.Chjoaygame (talk) 02:14, 5 June 2023 (UTC)
Attard, P., 2002, Thermodynamics and Statistical Mechanics: Equilibrium by Entropy Maximisation, Elsevier, Academic Press.
Many texts tell this story of friction. There is no mention here of microscopic factors such as unorganised motion of molecules. No mention of temperature here. Chjoaygame (talk) 19:09, 6 June 2023 (UTC)
Baierlein, R., 1999/2005, Thermal Physics, 6th printing, Cambridge University Press.
Evidently, for Baierlein, thermodynamic work is defined in terms of the external parameters of the system itself, without mention of ordinary physical work in the surroundings. For a closed system, for Bairlein, heating is defined by exclusion of thermodynamic work. Baierlein here forgets Joule's experiments.
Later, talking about the Carnot engine, Baierlein mentions rubbing and friction, but doesn't mention hammering or heat when he does so:
Conduction more or less implies that the source of the transferred heat has a temperature. Laser radiation is an example of heat transfer when the temperature of the source is not counted. Baierlein doesn't tell us whether frictional rubbing, viscous damping of fluid motion, or hammering, are counted as work or as heat.Chjoaygame (talk) 20:12, 6 June 2023 (UTC)
Blundell, S.J., Blundell, K.M., 2006, Concepts in Thermal Physics, Oxford University Press. Introductory comment:
Formal definition:
Heating by transfer of heat from one thermodynamic system to another, and heating by rubbing a thermodynamic system with something in the surroundings:
The conversion of ordinary physical work into heat occurs in the process of transfer.
In the section on the first law, some historical information. Lavoisier's 1789 notion of caloric. Thompson's 1798 heating by friction. Mayer's 1842 frictional generation of heat in paper pulp. Joule's frictional paddle experiment (1840 to 1845).Chjoaygame (talk) 06:14, 7 June 2023 (UTC)
Bridgman, P.W., 1943, The Nature of Thermodynamics, Harvard University Press.
Considering the first law, Bridgman analyzes heat production by friction:
I suppose the last phrase refers to the interface between the piston and the gas, not to that between the piston and the cylinder where there is friction. The difference between work done on the system, and work received by the system, by such a thing as a rotating paddle, was observed long ago by Bryan to be due to friction, and is the basis of the original experiments, by Davey, Thompson, Mayer, and Joule, that blew away the caloric theory and measured the mechanical equivalent of heat, but Bridgman is right to remark that this sort of thing does not occur very often in the conventional thermodynamic analysis. Perhaps the reason for this is that friction is not quite so easy to account for mathematically? Chjoaygame (talk) 06:41, 7 June 2023 (UTC)
Bryan, G.H., 1907, Thermodynamics: an Introductory Treatise dealing mainly with First Principles and their Direct Applications, B.G. Teubner, Leipzig.
Bryan was writing when thermodynamics had been established empirically, but people were still interested to specify its logical structure. The work of Carathéodory also belongs to this historical era. Bryan was a physicist while Carathéodory was a mathematician.
Bryan started his treatise with an introductory chapter on the notions of heat and of temperature. He gives an example of where the notion of heating as raising a body's temperature contradicts the notion of heating as imparting a quantity of heat to that body.
He defined an adiabatic transformation as one in which the body neither gains nor loses heat. This is not quite the same as defining an adiabatic transformation as one that occurs to a body enclosed by walls impermeable to radiation and conduction.
He recognized calorimetry as a way of measuring quantity of heat. He recognized water as having a temperature of maximum density. This makes water unsuitable as a thermometric substance around that temperature.
His second chapter started with the recognition of friction as a source of heat, by Benjamin Thompson, by Humphry Davy, by Robert Mayer, and by James Prescott Joule.
He stated the First Law of Thermodynamics, or Mayer–Joule Principle as follows:
He wrote:
He explained how the caloric theory of Lavoisier and Laplace made sense in terms of pure calorimetry.
Having rationally defined quantity of heat, he went on to consider the second law, including the Kelvin definition of absolute thermodynamic temperature.
In section 41, he wrote:
He then stated the principle of conservation of energy.
He then wrote:
On page 46, thinking of closed systems in thermal connection, he wrote:
On page 47, he wrote:
On page 48, he wrote:
I don't know whether Max Born knew of Bryan's work when he persuaded Carathéodory to undertake a mathematical investigation of the foundations of thermodynamics, or whether Carathéodory knew of Bryan's work, as he prepared his celebrated 1909 paper.
In my opinion, Bryan's definition of heat is the best available, and has been accepted by many thermodynamicists, and by the preponderance of past editors of this Wikpedia article. I regard Bryan's Treatise as a Wikipedia reliable source.Chjoaygame (talk) 04:48, 23 June 2023 (UTC)
Callen's first edition (1960), talking about closed systems (no transfer of matter), said on page 7:
Callen's second edition (1985, with material about statistical mechanics) is practically the same here.
No mention of temperature there. He was talking about an energy transfer into the system from its surroundings.
One may ask, did Callen admit that friction involves energy transfer to the hidden atomic modes? I presume that it does, and so that Callen's definition of heat includes heat production by friction, converting mechanical work into heat.
In the second edition, Callen continued:
Later, Callen continued:
Callen didn't say explicitly here whether the work is done with our without friction. But he continued by writing:
Callen assessed the work done as that done through that done in turning the crank. Callen regarded such work as measuring the difference in internal energies of the final and initial states of the working body.
He went on to write:
He didn't explicitly say that Joule's experiment did not measure a quantity of heat.
He went on to use the concept of quasi-static processes, indicating that he would later define them, using inexact differentials, and writing:
Thus his inexact differential is effectively translated into an exact differential
In this view of Callen, it seems that Joule did not measure a quantity of heat produced by friction. He measured a change in internal energy, and then equated that to a quantity of heat measured by calorimetry.
Soon, Callen wrote, on page 27:
Evidently, Callen regarded the work of Carathéodory as masterly. Callen didn't use differentials of heat to define entropy. Callen just postulated its existence as an extensive state variable.Chjoaygame (talk) 08:31, 11 July 2023 (UTC)
Çengel, Y.A., Boles, M.A., Kanoğlu, M., Thermodynamics: An Engineering Approach, 9th edition, 2019. Talking first about the zeroth law, on page 17:
In my opinion, this really belongs to the preliminary statement that thermodynamics deals with bodies in their own states of internal thermodynamic equilibrium, and in equilibrium with connected bodies. The authors then proceed to talk about temperature, not waiting for the second law to tell them how to define thermodynamic temperature.
On page 56, the authors write:
On page 60, they write:
On page 62, they write:
On page 65, they write:
They do not explain precisely how these electromotive forces act through a distance.
On page 66, they focus on mechanical work, writing:
They apparently feel no need to remark that their aforementioned electrical work violates their second requirement for "work": the boundary doesn't move. This is, however, a mere pedagogical convenience, or perhaps oversight, for they subsequently treat electrical separately from mechanical work.
On page 66, they write:
They do relate this to the distance moved by a point in the rotating parts, measuring distance moved in a suitable way. But the boundary of the system is not displaced and doesn't move.
No mention so far of friction, but on page 68, they write:
On page 70, they discuss other forms of work, including electrical work:
These engineers have discarded the foundational idea that friction generates heat. Thus, they reverse the thermodynamic tradition expounded by Bryan and by Carathéodory of defining heat by its not occurring through thermodynamic work, nor through transfer of matter. These engineers define thermodynamic work by its not occurring through heat.Chjoaygame (talk) 13:16, 22 June 2023 (UTC)
The reversal of the orthodox thermodynamic tradition is an example of circular reasoning. Quantity of energy transferred as heat is defined in terms of "work", while "work" is defined in terms of heat. This might be resolved by giving "work" two distinct and logically unconnected definitions, committing an act of logical equivocation. Wikipedia should avoid logical equivocation in its presentations.
More should be said. While it is possible to define empirical temperatures and so to define quantity of energy transferred as heat when it is between two bodies that each possess its respective temperature, it is convoluted and undesirable to need to redefine thermodynamic temperature after the acceptance of the second law. One might argue that the second law precedes the first law in logic, though that would be a hard road to hoe: how to define entropy without a prior definition of the distinction between heat and work transfers. A further problem with defining heat in terms of temperature difference is that not all sources of radiation, nor all sources of conducted heat, have a definite and uniquely defined temperature, so that the two-body definition of heat transfer doesn't work. From the orthodox thermodynamic viewpoint, the two-body temperature difference definition is merely a special case of the proper general definition by exclusion of thermodynamic work. The solution to the latter problem is to stay with the traditional thermodynamic base case of a system and its surroundings: the system has properly and uniquely defined thermodynamic state variables; but the surroundings are not so constrained; in the surroundings, from the viewpoint of thermodynamics, anything goes.
These points were, over several years long ago, established in this talk page, but they seem to have been discarded recently. The article should perhaps thoroughly explain this logic.
For Wikipedia, a good resolution is to avoid circular and convoluted reasoning, and to stay with the orthodox thermodynamic position, and to declare, solely for this specific purpose, that Çengel, Boles, and Kanoğlu 9th edition is not a reliable source.Chjoaygame (talk) 03:16, 23 June 2023 (UTC)
Denbigh K., The Principles of Chemical Equilibrium: with Applications in Chemistry and Chemical Engineering, 4th edition, 1981, Cambridge University Press.
On page 10, Denbigh wrote:
This is thinking in terms of calorimetry, not so much in terms of the idea of Locke, Thompson, Mayer, and Joule, that friction generates heat.
On page 18, Denbigh wrote:
Denbigh is thinking of heat as defined by calorimetry, relying on two bodies, each with a defined temperature. This thinking seems to show that, after all, Joule did not measure the mechanical equivalent heat, because he did not heat the water in the vat. Denbigh would perhaps reply that, nevertheless, the water changed its internal energy by the same amount as it would change if the process had been one of heating by conduction and radiation. That reply, however, relies on the Joule experiment having converted all the externally applied work into heat. If only some of the externally applied work were transformed into heat, more complicated reasoning would be necessary, as indicated by Bryant and by Bridgman. This may explain why Planck talks of surface rubbing as distinct from internal friction.
Denbigh on page 19 wrote:
This might be called 'the principle of calorimetry', and is the main basis of the caloric theory of heat. Bryant and Bridgman would reply that they do not accept such an equality as a general principle. It does not cover friction in such processes as drilling, rubbing, grinding, or hammering.
Denbigh made some concessions to the idea of thermodynamic work as defined by changes of state variables other than thermal, i.e., other than entropy and temperature. For discussion of heat, the internal energy is always one such state variable. Denbigh mentioned Joule's measurements of the energy changes due to friction between iron blocks, but he did not elaborate. He also wrote the following:
For example, in hammering. This didn't deal with rubbing. Apart from the latter two brief mentions, Denbigh conveniently ruled out discussion of friction.
For the definition of work, however, Denbigh wrote:
The 'work' terms are sometimes taken to exclude 'chemical work'.
Eventually, Denbigh disallows the heat produced by friction. He relies on a roundabout calorimetric definition, supposing that heat comes from a convenient "heat bath" of water at its temperature of maximum density, and writing:
A chemist is hardly interested in friction at this stage of thinking, and this perhaps explains why Denbigh defies Thompson, Mayer, and Joule, and rules out friction as a generator of heat. Friction is more the province of physicists such as Planck.Chjoaygame (talk) 11:46, 1 July 2023 (UTC)
Dugdale wrote on page 4:
On page 20, he wrote:
On page 21, he wrote:
Dugdale gives plenty of detail on Thompson's and on Joule's experiments, described by them as production of heat by friction, but he interprets them as occurring through work. Nevertheless, he gives much attention to calorimetry, as noted above.
On page 21, he wrote:
Dugdale is not interested in friction, and puts his faith in a principle of calorimetry, that the heat lost by one body is gained by the other. He defines work in terms of what happens in the surroundings, apparently not in terms of the change it produces in the state variables of the system. Yet, for the interpretation of experiments, he also relies on the definition of the state of the system in terms of 'work' variables, including internal energy.Chjoaygame (talk) 10:24, 30 June 2023 (UTC)
Dunning-Davies, J., 2011. Concise Thermodynamics: Principles and Applications in Physical Science and Engineering, 2nd edition, Woodhead Publishing, Oxford UK.
On page ix, Dunning-Davies wrote:
Evidently, Dunning-Davies is not too concerned or overfamiliar with the cave man's ability to generate heat by friction between sticks, or with the coachman's concern that sometimes the heat generated by friction of his axle with it bearing can set the coach on fire, or with the blacksmith's heating of his work by hammering. Those primitive fellows are not overfamiliar with the caloric theory of heat. Apparently, Dunning-Davies is happy with the ordinary language word 'heat' as a scientific term. He sees no need to explicitly define it in this context. He will remain content to talk of two systems being brought into thermal contact. Does that include mutual radiative exposure?
On page 1, Dunning-Davies, without explicitly defining it, introduces a term 'thermal properties' as follows: he remarks that "...the laws of thermodynamics ... are simply expressions of common experience of the thermal properties of matter and radiation." On page 2, again without explicitly defining it, in the context of the caloric theory, he uses the term "thermal contact".
On page 5, Dunning-Davies focuses on two bodies that each possess a temperature, not worrying about the general approach to thermodynamics that requires the system to have a temperature, but does not impose any such requirements on the surroundings. He wrote:
On page 13, Dunning-Davies wrote:
Dunning-Davies is not precise about exactly how to specify adiabatic and non-adiabatic work.
On page 26, Dunning-Davies wrote:
Evidently, Dunning-Davies knows the highly mathematical work of Carathéodory but seems unimpressed by the prior more thermodynamic work of Bryant.
Dunning-Davies so loves the caloric theory of heat that he sets up the above roundabout way of defining heat through mechanical processes only, apparently not impressed by Thompson's, Mayer's, and Joule's methods of allowing straightforward mechanical definition of quantity of heat by measuring its production in friction. Why do things the easy and obvious way when you can do them a hard way?Chjoaygame (talk) 10:06, 1 July 2023 (UTC)
Giles, R., 1964, Mathematical Foundations of Thermodynamics, Pergamon Press, Oxford UK.
On page 1, Giles wrote:
On page 2, introducing another way of defining entropy, Giles wrote:
Still on page 2, about a better way of defining entropy, Giles wrote:
After some consideration of frictionless mechanical processes, and then explicitly defining entropy, looking at a thermodynamic system A connected to a mechanical device M in the surroundings, on page 109, Giles wrote:
On page 115, Giles wrote:
As Giles observes, his definition of heat here relies on the concept of a quasi-static process. Is this perhaps necessary for a definition of heat? For example, the definition of thermodynamic work, for a finite increment of volume at constant pressure requires that the process be slow enough to allow the pressure of the system to be defined throughout it; this definition also requires that the process be slow enough to allow the temperature to be defined throughout it. This definition demands the simultaneous definition of heat and thermodynamic work.
We may observe that this definition of heat does not consider such an abrupt process as hammering to produce heat. Hammering, drilling, fluid friction, and rubbing, convert energy from the surroundings directly into heat. The energy from the surroundings can be measured directly without regard to the intimate details of the process, and, provided it is all converted to heat, it can measure the quantity of heat directly. That is the merit of the works of Thompson, Mayer, and Joule. It might reasonably be objected that in such processes, some of the energy from the surroundings is converted into heat in the surroundings. For example, the hammer will become hot. It seems to follow that some kind of idealization is necessary for the precise definition of thermodynamic quantities.
Eventually, Giles defines heat through his prior definition of entropy. He relies on entropy as measuring the whole irreversibility of any thermodynamic process. He is presenting the idea of heat after he has settled on the second law of thermodynamics. This is reasonable and logically defensible, though it is not the commonest way to define heat. This mathematically oriented reasoning of Giles competes with the older physically oriented work of Clausius, that defined entropy in terms of infinitesimal increments of heat, and with the older physically oriented work of Bryan and mathematically oriented work of Carathéodory that defined heat as a residual from work. It is often felt that defining things in terms of adiabatic work is straightforward, and is evidently based on simple physics.Chjoaygame (talk) 09:41, 3 July 2023 (UTC)
Grandy, W.T., 2008, Entropy and the Time Evolution of Macroscopic Systems, Oxford University Press, Oxford UK.
Leading up to an account of entropy, on page 3, Grandy wrote:
Grandy does not go on to consider the details of thermodynamics, such as an explicit definition of 'heat', because he is concerned with statistical mechanics.Chjoaygame (talk) 05:08, 4 July 2023 (UTC)
Guggenheim, E.A., 1967, Thermodynamics: An Advanced Treatment for Chemists and Physicists, North Holland, Amsterdam.
On page 9, Guggenheim excluded friction from his book, writing:
On pages 9–10, Guggenheim wrote:
On page 12, Guggenheim wrote:
Guggenheim is clear that heat cannot be generated by friction. The caveman, the coachman, and poor old Thompson, Mayer, and Joule would turn in their graves. One wonders what did happen in the experiments of the latter authors. Work was converted into 'energy'. Into internal energy? Work is a process term; internal energy is a state term. Heat is a process term.
Evidently, Guggenheim recognizes that, in thermodynamics, heat is a process notion, and that it is defined by exclusion of work.
But also evidently, Guggenheim isn't interested in the tradition that thermodynamics is based on statements about a system and its surroundings, while he prefers to think in terms of two interacting thermodynamic systems; and, not admitting the idea of friction, he isn't interested in the Thompson–Mayer–Joule–Bryant–Bridgman–Planck idea that it generates heat, so that it isn't necessary that and . Friction is essentially a process notion, referring to the surroundings, and is not discussed in detail in thermodynamics; only its effects are recognised there.Chjoaygame (talk) 07:04, 4 July 2023 (UTC)
Keenan J.H., 1941, Thermodynamics, John Wiley & Sons.
On page 6, Keenan wrote:
On page 9, referring to paddlewheel experiments without mentioning Joule's name, Keenan wrote:
No mention there of friction. Actually, Keenan here inverts the usual idea that the 'system' is the calorimetric vat, and that the 'surroundings' are the location of the falling weight.
On page 67, omitting from his book mention of Joule's measurements of the mechanical equivalent of heat, and having excluded friction till this point, Keenan wrote:
On page 114, Keenan observed that the work done by the prime mover may exceed the work received by the moved element, the difference being due to friction:
This is in accord with the view of Bryan and Bridgman, but Keenan does not enter it into his definition of heat. Keenan neatly avoids talk of heat here, just talking about "work absorbed". What is 'absorption' of work?
Likewise, on page 132, Keenan avoided talk of heat when he wrote:
"Losses of work"? What does that mean? Keenan wrote anything except that friction produces heat.Chjoaygame (talk) 11:22, 4 July 2023 (UTC)
Kittel, C., Kroemer, H. (1980). Thermal physics, 2nd ed., W.H. Freeman, USA.
On page 227, Kittel & Kroemer wrote:
The authors seem to assume that, for a closed system, transfer of energy is either by work or by thermal contact with a reservoir. Yet, for them, work is the transfer of energy to a system by a change in the system's external parameters. Here, as external parameters, they list volume, magnetic field, electric field, and gravitational potential. The present note considers volume to be a prototypical external parameter, but magnetic and electric fields are not so simple. At this point, Kittel & Kroemer mention electric field, an intensive variable, but they say nothing about the dielectric polarization that is its conjugate extensive state variable. They do not say precisely what they mean by an 'external parameter'. They seem to partly exclude friction from their considerations. But on page 232, they wrote:
For conversion of heat into work, it is evident that they refer to devices that operate through a sequence of several processes each initiated by respective thermodynamic operations, and that they are not thinking of a single separate process initiated by a single separate thermodynamic operation. On the other hand, it does seem that friction can be considered as a single separate process that converts work into heat.
Nolting (2017) defines 'work' by listing the various forms of thermodynamic work, defined by integrals of conjugate thermodynamic state variables. He makes no particular reference to external mechanical force × distance notions.
On page 29 he writes about heat:
This says nothing about latent heat, and it is a pity that it doesn't mention internal energy at this stage. Consequently, heat is here based partly on temperature, which has not yet received its thermodynamic definition.
Nolting goes on to write:
It seems that, instead of our preferred idea of defining entropy through heat, Nolting is happy to define heat through entropy, though he hasn't yet presnted the second law. Here, he goes on to define temperature through entropy. It seems that Nolting is not too fussed about logical development through empirical facts.
In his Textbook of Thermodynamics (1913), Partington writes:
Partington also wrote:
Partington thinks that heat and work are interconvertible, and that Joule proved it:
He further writes:
So Joule was indeed measuring the mechanical equivalent of heat.
According to Petrucci, Herring, Madura, and Bissonnette, 11th edition:
They also show a picture of Joule, with the caption:
Evidently, Joule did measure quantities of heat, but it remains perhaps mysterious how he did so, if the above definition is adhered to. We may remark that there, the authors refer to the law of conservation of energy, not to the first law of thermodynamics.
On page 14, Pippard writes:
In effect, Pippard is rejecting the relevant definition of work as 'thermodynamic work' defined by the change of thermodynamical state variables. He is defining work as defined in the surroundings, mechanically, without regard to its being an ingredient of a thermodynamic process.
Pippard then writes:
Considering Joule's paddle wheel experiment: How do we define 'the system'. Pippard seems to think that the paddles are part of the system, and are adiabatically separated from the surroundings. He doesn't actually demand measurement of the change of volume of the water. It may tacitly be assumed that it is zero, I think?
Why does Pippard think that Joule measured the mechanical equivalent of heat, when he thinks that the experiment did not add heat to the water? Perhaps he doesn't really believe that Joule did measure the mechanical equivalent of heat? He writes:
After all, Pippard seems to think that Joule didn't measure the mechanical equivalent of heat. Oh, dear! All of Pippard's care for logical rigour has overthrown Joule !! What a clever fellow is Pippard !! Planck was wasting our time rattling on about friction !!
Pippard wriggles out of this by writing:
How does Pippard eventually define heat? He does so partly by dismissing the thought of the paddle being considered as part of the surroundings, and as itself being heated by the performance of mechanical work. He requires perfectly adiathermal performance of mechanical work:
He writes:
with heat. In point of fact the properties concerned are not many in number, and may be summarized as follows: (1) The addition of heat to a body changes its state. (2) Heat may be conveyed from one body to another by conduction, convection or radiation. (3) In a calorimetric experiment by the method of mixtures, or any equivalent experiment, heat is conserved if the experimental bodies are adiabatically enclosed.
Evidently, Pippard thinks that Maxwell was mistaken to say that convection is transport of internal energy, not heat transfer as such. He doesn't admit that heat can be generated by friction. How does convection transfer heat from one body to another? My impression is that, according to Maxwell, convection occurs within a fluid body during a process, and is not a transfer of energy as heat from one body to another.
I would say that Pippard defines quantity of heat calorimetrically, especially when I see him write that "heat is conserved if the experimental bodies are adiabatically enclosed". But if the bodies are not adiabatically enclosed, heat is transferred by conduction or radiation.
So this is a stumbling block for the definition of heat as transfer of energy other than by thermodynamic work or the transfer of matter. It is the calorimetric definition of heat. Evidently, according to Pippard, mechanical work cannot be transformed into heat. Poor old Joule.
Reif, F. (1965). Fundamentals of statistical and thermal physics, McGraw-Hill, New York.
This standard text is oriented by statistical mechanics. The statistical mechanical world is not that of macroscopic thermodynamics. Statistical mechanics is not interested in general in the system's surroundings, where hammers, pulleys, and frictional and other devices are located. For statistical mechanics, more interest lies in such interactions as between two bodies each initially in its respective thermodynamic equilibrium.
The relevant point here is that work is defined by changes in the system's external parameters, not by the arrangements of frictional and other devices in the surroundings.
On page 68, considering processes in closed systems, Reif wrote:
By 'external parameter', Reif means extensive thermodynamic state variables other than the thermodynamic variables entropy and internal energy. The prime relevant example of such an 'external parameter' is volume. Such variables can be measured externally to the working body, for example by measuring the length, breadth, and width of the working body. Reif means to exclude internal variables such as the various intensive state variables such as pressure and temperature, and to exclude transient density fluctuations within the body; the latter are supposed to average out to zero over time in thermodynamic equilibrium.
When all the external parameters remain unchanged in a process, the internal parameters can still change. It is these that change in the kinds of process that Reif is considering. For example, temperature and pressure can change in such processes. For example, when the body is heated, the pressure and temperature will usually increase. In bodies composed of some mildly unusual substances, such as water at temperatures in a certain range near freezing, heating with increase in temperature can result in decrease of pressure. (This makes water unsuitable as a thermometric substance in that temperature range.)
The situation can be described in more detail in the light of the second law of thermodynamics, which defines entropy and thermodynamic temperature
If the process, defined by different initial and final external parameter values, is imagined as isentropic, devoid of friction, and adiabatic (no heat conduction, no radiative transfer), then the quantity of work done by the system is mathematically uniquely defined, and is given by an integral of the form
that describes any such process obeying the fundamental equation of the system throughout.
If, on the other hand, the process is imagined as entirely due to heat transfer, in Reif's words, "purely thermal", then an integral of the form
will describe any such process obeying the fundamental equation of the system throughout.
One can consider more general processes, irreversible, with both work and heat transfer, including friction, obeying the fundamental equation of the system throughout, for which
Such integrals are path integrals, for which the path is not uniquely determined. This is why the symbols and are said to denote inexact differentials.
In this view, Joule did not do thermodynamic work on his vat of water; instead, he heated it through friction. So he directly measured the mechanical equivalent of heat, a substantial achievement, commemorated in the name of the SI unit of quantity of energy.
Reif then specializes to consider a case in which the surroundings are of the same character as the working body. This excludes such processes as stirring with a paddle, because the paddle is driven by devices in the surroundings that are quite different in character from the working body. Such processes are not the main interest of statistical mechanics, but they belong to thermodynamics in general.
Callen's first edition (1960), talking about closed systems (no transfer of matter), said on page 7:
This is the converse of Reif's definition of heat, in the sense that Callen's "hidden atomic modes" are the subject of Reif's "first kind of interaction where the external parameters of the system remain unchanged".
Schroeder, D.V., Introduction to Thermal Physics, 2000, Addison Wesley Longman, San Francisco. Schroeder is writing from the viewpoint that starts from microscopic physics, as against the thermodynamic viewpoint that starts with macroscopic physics. He leads with
Schroeder does make the important point that "both heat and work refer to energy in transit.
But, like Çengel et al., Schroeder defines thermodynamic work by exclusion of heat.
This reverses the logic of Bryant and Carathéodory, who define heat by exclusion of work. Schroeder's definition of work, by exclusion of heat, is in contrast with the usual idea in physics, that work is defined mechanically, in terms of such things as the ability to lift a weight. Thermodynamics began with the idea that spontaneous work is done by the working body of a heat engine, not with Schroeder's idea that
Schroeder doesn't seem to consider friction within the system such as occurs in Joule's paddle wheel experiment.
In the above quote from Schroeder, there is no mention of friction. His first mention of friction is in the following footnote:
Though Schroeder is not focusing on friction, that remark is compatible with the above remark of Bryant:
And with the above remark of Bridgman:
This may be interpreted by saying that ordinary physical work done by a mechanism in the surroundings must sometimes be distinguished from thermodynamic work done on its surroundings by a thermodynamic system, as for example by a heat engine.
Perhaps it is, as Schroeder thinks, "strange to think that there is no "heat" entering your hands when you rub them together to warm them up." Perhaps rubbing one's hands together mainly has the effect of increasing the blood flow through the hands? It is not quite the same as rubbing two pieces of ice together.
Classical theoretical texts on thermodynamics define changes in the internal energy of a thermodynamic system strictly in terms of thermodynamic work done by a body enclosed in a container with adiabatic walls. In practice, perhaps most measurements of change in internal energy are done by calorimetry, as for example in Joule's paddlewheel experiment.
In thinking about thermodynamic work, one should bear in mind that thermodynamics is primarily about differences between thermodynamic states. This is why thermodynamic work is defined by differences between thermodynamic states. Thermodynamics is not simply about forces that an "agent" in the surroundings can exert to do work on the thermodynamic system. It is about forces that a thermodynamic system can exert to do work on its surroundings; such work can be received in the surroundings partly as work against friction, i.e., as heat.
It is perhaps worth remarking at this point that "chemical work", referring to such quantities as , might safely be called 'chemical work-like change'. This is because "chemical work" is defined neither by mechanical forces that the surroundings exert on the system, nor by mechanical forces that the system exerts on the surroundings, but by changes in the state variables of the system.
Schroeder is a chemist who approaches thermodynamics as secondary to microscopic physics, apparently not a physicist who learnt thermodynamics as a macroscopic topic from Carnot, Joule, Mayer, Joule, Bryant, Carathéodory, and Planck. One may ask, which is better for Wikipedia, that it give priority to the thinking such as Fourier's, in terms of partial differential equations and the caloric theory, or that it reflect the knowledge of a cave man or pre-industrial coachman, that friction generates heat? I am inclined to bear in mind that Wikipedia is often enough quoted just from the first defining sentence of an article, as if that authoritatively settles a question.Chjoaygame (talk) 04:09, 29 June 2023 (UTC)
Zemansky, M.W., Dittman, R.H. (1997). Heat and Thermodynamics: An Intermediate Textbook, 7th ed., McGraw-Hill, New York.
Zemansky & Dittman
On page 49, defining thermodynamic work, the authors wrote:
On page 73, the authors wrote:
On page 78, the authors wrote:
On page 80, the authors wrote:
On page 81, the authors wrote:
The authors' consideration of heat and the first law is only partly done at this stage. They consider it further later in their text. They continue here by writing
The authors are defining 'work' only in terms of generalized force and generalized displacement as measured for the surroundings, not in terms of such quantities as measured for the working body itself. Their result is that 'work' as measured by their definition leads to two mathematical changes in the working body, namely 'work' as measured for the working body itself by an integral in terms of the system's internal variables, and a quantity defined in terms of two characteristically thermodynamic variables, , the thermodynamic temperature, and , the differential of the entropy which are defined only through the second law. Thus, in effect, they might define two kinds of work, that defined for the working body, and that defined for the surroundings. They do not talk about this distinction, and so they hold fast to the caloric definition of 'heat', in contrast to that apparently assumed by Thompson, Mayer, and Joule, in terms of mechanics, by which work coming from the surroundings is converted by friction to heat going into the system. Some textbook writers are careful to notice this distinction: for example, Bryan, and Bridgman, who recognize the two kinds of 'work'. On the no-distinction side, for example, Guggenheim (1967) recognized only one kind of 'work' and ridiculed the idea of surroundings work being converted by friction into system heat; on page 12, he wrote
Authors who work with the no-distinction idea do the integrations for the calculation of work in a process indirectly, by integrating increments of heat measured by calorimetry. Apparently they are not interested in the integrations for 'work' in terms of the state variables of the system. The result is that 'work' done on the system by the surroundings is always in physical reality contaminated by 'heat' in a way that is not evident or directly traceable in the quantity of 'work' as defined by the processes entirely within the surroundings. In other words, for them, the differentials and do not measure 'heat' directly. The caloric theory of heat lives on.
Borgnakke, C., Sonntag, R.E. (2009). Fundamentals of Thermodynamics, 7th edition, John Wiley & Sons, Inc.
On page 90, they wrote
Very often, thermodynamics assumes that ordinary physical work is well defined, and so it takes 'work' as the basic quantity of energy, and then defines 'heat' as a residual from 'work', as do Borgnakke & Sonntag. (Some writers, e.g. Schroeder (2000), and perhaps Çengel, Boles & Kanoğlu (2019), define heat by calorimetry and then define work as a residual from heat.)
On page 90, they also wrote
The logic of definition here is perhaps thrown into doubt by the following sentence, which they wrote on page 104:
The logic of Borgnakke & Sonntag seems to require that the work done measured by the lifting of the weight should be equal to the work measured by the integral.
In the first above sentence, we are talking about the putative 'sole effect' of thermodynamic process; we have not heard explicitly about the thermodynamic operation that initiated the process, or how it is ensured that the process be quasi-equilibrium, or how concomitant heat transfer is ruled out. The second sentence, about "the integral of the product of an intensive property and the change of an extensive property", clarifies things. Physically thinking, concomitant heat transfer can partly be ruled out by requiring that the working body be contained by adiabatic walls. If we think, for example, of a gas expanding against the resistance of a piston in a cylinder in the surroundings, we could require the piston to be placed so that it can move up only by lifting a weight on top of it. We must think of the friction of the piston against the cylinder. If it is zero because of excellent lubrication, the gas itself will determine the speed of the lift. If the weight is great, it will move upwards slowly. A requirement is that the motion must be slow enough that the pressure in the gas remains defined throughout the process. This is necessary for the correctness of the above integral requirement. Do we require also that the upwards movement should not transiently overshoot its final position? Yes, we do require that. If the gas makes itself oscillate about its final position, then there will be internal friction, within the gas, that diverts some of the energy of expansion to remain as internal energy in the gas, so that the whole of the energy of expansion did not eventually reach the weight to lift it. Some of the energy of expansion is dissipated as heat in the gas. The process was not devoid of heat transfer and it was not isentropic; there was some heat transfer from the surroundings to the gas, when the weight gave up some of its potential energy as it compressed the gas. We do not have throughout the process that .
If the raising of the weight is to be isentropic, then the pressure–volume relation should measure the work done. There is no auxiliary variable to allow moderation according to the principle of Le Chatelier. But if the process is not isentropic, then the possibility of entropy change will allow moderating feedback through temperature change.
Dated 22:48, 8 August 2001, edit https://en.wikipedia.org/w/index.php?title=Heat&oldid=256829.
Comment: Heat is like a substance that flows. No mention of friction. Considers convection as a form of heat transfer.
Dated 15:50, 15 December 2005, edit https://en.wikipedia.org/w/index.php?title=Heat&diff=prev&oldid=31477735.
Comment: a new approach, immediately undone.
Dated 09:36, 23 January 2006, edit https://en.wikipedia.org/w/index.php?title=Heat&diff=next&oldid=36281983.
Comment: No mention of friction. Muddled thinking: "can be created by chemical reactions"; I would say that chemical potential energy can convert into internal energy; internal energy can be transferred as heat.
Dated 08:42, 25 January 2006, https://en.wikipedia.org/w/index.php?title=Heat&diff=next&oldid=36538164.
Comment: mentions creation of heat by friction.
Dated 01:54, 3 March 2006, https://en.wikipedia.org/w/index.php?title=Heat&diff=next&oldid=41816652.
Comment: using the phrase "energy in transit".
In ordinary English, heat is not a process: it is a quantity of energy. In ordinary English, heating is a process. In physics, heat is energy in transfer by certain mechanisms. Often enough, in ordinary English, and in eighteenth century physics, with the caloric theory, heat is perhaps vaguely a substance. But we do not defer to the caloric theory, because, by the nineteenth century in physics, the work of Rumford, Mayer, Joule, and others had shown that heat is generated by friction, thereby replacing the caloric theory with the mechanical theory of heat. For thermodynamics, friction is the conversion of energy of mechanical motion of contiguous bodies into heat. For thermodynamics, this may be explained, though not defined, by saying that heat is macroscopic energy in transfer through atomic, photonic, or other microscopic mechanisms. Callen writes "But it is equally possible to transfer energy via the hidden atomic modes of motion as well as via those that happen to be macroscopically observable." [1]
This book uses the word 'matter' as ordinary language, but, in near, though not exact, agreement with the poster of the edit, it once uses the phrase 'transfer of mass'. It writes neither 'mass transfer', nor 'transfer of matter'.
Page 222:
Page 30:
Page 121:
Page 131:
Page 157:
Page 270:
This book uses the word 'matter' as ordinary language. It uses the phrases 'matter transfer' and 'transfer of matter', but does not use the phrase 'mass transfer'.
Page 5:
Page 17:
Page 26:
Page 58:
This book uses the word 'matter' as ordinary language. It uses the phrases 'transfer of matter', and 'transfer of material', but not the phrase 'mass transfer'.
Page 5:
Page 6:
Page 77:
Page 81:
This book does not use the phrase 'mass transfer' where the poster of the edits would expect it. This book uses the word 'matter' as in ordinary language, and the phrase 'flow of matter' in a place where the poster of the edit would expect 'mass transfer'.
Page 51:
Page 59:
Page 157:
Page 327:
Page 328:
Page 338:
Page 341:
Page 342:
Page 343:
Page 345:
Page 348:
Page 366:
This book does not use the phrase 'mass transfer' where the poster of the edits would expect it. This book uses the word 'matter' as in ordinary language, and the phrases or clauses 'passage of matter', 'exchange of matter', 'matter exchange', 'it cannot exchange matter', 'no matter could be transferred', 'addition of matter to the phase from without', 'flow of matter', and 'transfer of matter', in places where the poster of the edit would expect 'mass transfer'.
Page 7:
Page 9:
Page 13:
Page 16:
Page 18:
Page 51:
Page 108:
Page 117:
Page 119:
Page 121:
Page 129:
Page 145:
Page 155:
Page 203:
Anderson (2005) 2nd edition page 8: "To learn more about chemical reactions, we have to become a bit more precise in our terminology and introduce some new concepts. In this chapter, we will define certain kinds of systems, because we need to be careful about what kinds of matter and energy transfers we are talking about; equilibrium states, the beginning and ending states for processes; state variables, the properties of systems that change during reactions; processes, the reactions themselves; and phases, the different types of matter within the systems."
Anderson page 10: "This is because we must be able to control (conceptually) the flow of matter and energy into and out of these systems." "Isolated systems have walls or boundaries that are rigid (thus not permitting transfer of mechanical energy), perfectly insulating (thus preventing the flow of heat), and impermeable to matter."
Anderson page 11: "Isolated system. Nothing can enter or leave the system (no energy, no matter). Whatever is inside the walls (which could be anything) will have a constant energy content and a constant composition. (b) Closed system. The closure is a piston to indicate that the pressure on the system is under our control. Energy can enter and leave the system, but matter cannot. The system here is shown as part liquid, part gas or vapor, but it could be anything. Both the liquid and the gas could also be considered as open systems, inside the closed system. Each may change composition, although the two together will have a constant composition. (c) Open system. Both matter and energy may enter and leave the system. The system may have a changing energy content and/or a changing composition. The pitcher shows one way of adding matter to the system."
Çengel & Boles (2019) 9th edition page 12: "Matter is made up of atoms that are widely spaced in the gas phase. Yet it is very convenient to disregard the atomic nature of a substance and view it as a continuous, homogeneous matter with no holes, that is, a continuum."
Seamless Wikipedia browsing. On steroids.
Every time you click a link to Wikipedia, Wiktionary or Wikiquote in your browser's search results, it will show the modern Wikiwand interface.
Wikiwand extension is a five stars, simple, with minimum permission required to keep your browsing private, safe and transparent.