Talk:Heat/Archive 3

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What is this page for?

It seems to me that a more enlightening question for a straw poll would be to ask, what is this page for. It seems to me that there are at least two different lines of thought:

1. Heat only in a narrow first law of thermodynamics sense: δQ only.
2. Heat in a broader sense, eg "a name for internal energy when we are thinking of the energy as randomised and free to move from one part of the system to another or between systems...," as the textbook by Waldram (1985) quoted above puts it.

Most of the historical material recently put back into this article by Sadi uses the word heat in the second (broader) sense, not the first (narrow) sense. This is also the way "heat" tends to be presented at an elementary level to ten-year-olds onwards, with no distinction made between heat and thermal energy; with the narrower thermodynamic definition left until college teaching.

So the question I think would be useful to canvas opinions on by way of a straw poll, is what should be the scope of this article? The narrow first-law use only? Or all uses of the term heat, both now and historically too? -- Jheald 14:16, 2 July 2007 (UTC)

Of course the article is intended to be written in the broader sense, i.e. all perspectives. You are welcome to add material to make it so. If you are going to make efforts to further distinguish between heat and thermal energy, however, be sure to use support references that explicitly use the word "thermal energy" in contrast to “heat”. Thanks: --Sadi Carnot 15:14, 2 July 2007 (UTC)

Sadi, I don't think there is any "Of course" about it. My impression from comments here is that most editors want this article to be distinctively about (1), not (2). But to see just where the balance of consensus lies, that's why I think this question is actually quite a useful one for a straw poll. Jheald 15:49, 2 July 2007 (UTC)

Jheald, at this point, I really don’t know what opinion you are trying to poll. If it has to do with moving the entire heat page to heat (thermodynamics), that idea has been voted down. If, however, you want to add new material in Wikipedia (somewhere) focused on #1 or #2, I don’t think that anyone will object. --Sadi Carnot 16:18, 2 July 2007 (UTC)

Sadi, I don't know how to put this more straightforwardly, but the contention I am trying to explore here is: that material which only relates to sense #2 and not sense #1 should not be featured in this article and should be removed.

If you disagree, I suggest you register that disagreement by supporting "All uses of the term heat" below as the appropriate scope for the article. Jheald 16:28, 2 July 2007 (UTC)

You seem to be speaking in riddles, but I assume that you want to move the history “section” (that I wrote) out of the article? Maybe, in the future I would support a partition out of some material to history of heat when the article gets longer; presently, however, the article is only 31kb and I really don’t see a need for this (maybe when the article goes to 45kb or more). Consensus has clearly shown that people are against merging anything. If there is something else that I am missing in your view, then please clearly say: “I want to move this section here or there, etc…” Thanks: --Sadi Carnot 17:49, 2 July 2007 (UTC)

Narrow first-law use of the word only

1. Support. In my view the δQ use is sufficiently important, sufficiently precise, and students find it sufficiently difficult, that this article deserves to focussed on it solely. Material about the use of the word heat in a less narrow, more general, elementary, and/or historical sense I think should be in a separate article, probably thermal energy. Each article should mention the other usage of the word heat only in passing, to qualify that that is what the article is not about. To underline the point, I would support titling this article heat (thermodynamics) to emphasise it being about the narrow, first law meaning of the word, not the broader/introductory/more general/less precise meaning. Jheald 14:16, 2 July 2007 (UTC)

Moving the entire heat article to heat (thermodynamics) was obviously a wrong move, and was reverted back by User:Hoary here. If, however, you want to add new material, via a new header in this article, on how quantities of heat "δQ" are used in the laws of thermodynamics or in thermodynamics in general, I'm sure that would be useful (then possibly later when the article grows too long we sub-section it out?). The idea of the term heat (thermodynamics) is redundant (as pointed out by User:ScienceApologist). In the history of science, the phenomenon of heat evolved from mythological theories of fire, to heat, to terra pinguis, phlogiston, to fire air, to caloric, to the theory of heat, to the mechanical equivalent of heat, to thermo-dynamics (sometimes called energetics) to thermodynamics. Hence, any article on heat is thermodynamic. --Sadi Carnot 15:14, 2 July 2007 (UTC)

1. Support I would like to point out that to those who do not support narrow first-law use that the "broader" colloquial usage is actually much closer to first-law use than you may realize. When someone touches something and they say that it is hot they are reacting to the rate of thermal energy transfer not temperature differentials or total thermal energy of the system. Hot or coldness that we sense has to do when energy transfer rates alone. Degrees of freedom and heat capacities as well as effective surface area and thermal contact effect this and can change the rate of transfer at the same temperature differential (e.g. metal vs. plastic). We as human beings are only capable of detecting heat (energy transfer). And the use of heat as thermal energy in a physics sense is just 100% wrong.--Nick Y. 21:08, 2 July 2007 (UTC)

Nick, to clarify what Jheald is trying to do here (as far as I have gathered), is to gain support, via some ulterior motive first law platform, to move the history of heat section to thermal energy; a move that I have already reverted. As far as I have been able to get out of him, he doesn't plan on writing any new material, but only to move the history section. In any event, I wrote the history of heat (not the history of thermal energy), and will strongly oppose any move of the history section. If he wants to write a new history section, focused on the laws of thermodynamics in relation to heat, that’s fine; but a move of the history of heat anywhere other than heat or history of heat (a new article) is clearly not correct. --Sadi Carnot 02:33, 3 July 2007 (UTC)

I do not agree with moving the history of heat section. I do agree that heat is the transfer of thermal energy between two bodies. In my understanding it is an energy flux and is not energy. I also agree that there is a place for the historical usages, even of magical nature somewhere in wikipedia but this article should primarily focus on the modern technical definition of heat. --Nick Y. 19:33, 3 July 2007 (UTC)

There's no reason why it should be wrong. Rather, some people have gotten together and decided that it should be a different terminology than the colloquial idea. Thus we have the ridiculous state of afairs that objects have heat capacity, but never do contain heat. Or rather, they don't contain heat if it's evenly distributed-- but while it's in the act of distribution within the object, the part which is moving remains heat (subject to the heat transfer equations), but when it stops and temperatures equilibrate, this energy stops being heat and (voila!) becomes thermal energy. This is a head scratcher in terminology (as though we had two different terms for water when it was running vs still), and is going to need to be very carefully explained in the LEADS of all these articles, ESPECIALLY if "heat" is not going to be renamed "heat (thermodynamics)," and yet is no longer to be used in the old sense that it was used, when the term "heat content" was coined (and if truth were told, the people who made up the term heat capacity probably were thinking in terms of heat staying heat in that case as well, not magically turning from heat into thermal energy as soon as it got settled-- but nevermind). In either case, I don't care what you do with the aticles, so long as the LEAD explains that "heat" as we're now going to use it, means "δQ" and (perish the thought) not Q.

And by the way, what the nervous system senses is not heat transfer per se, but simple absolute temperature (and also dT/dt, the temp rate change). These are closely connected to dQ of course, but heat capacities on both sides of skin can change, and when they do, it's apparent that it's not heat, but heat's effect on tissue (temperature) that is what is felt. SBHarris 02:23, 3 July 2007 (UTC)

The heat capacity of our skin is pretty much a constant. Different material of the same temperature but different heat capacities feel very different to us because of the rate of heat transfer that, yes of course, changes the temperature. Heat and temperature are closely related but temperature differentials are not something we can accurately sense due to different heat capacities resulting in different rates of energy transfer (heat) which yes result in increased temperature. Once we are in equilibrium we may come to understand the temperature of the object more accurately, but that is temperature not a temperature differential. In any case my point was that many colloquial uses of heat are consistent with the modern scientific definition. Some may not be.--Nick Y. 19:33, 3 July 2007 (UTC)

All uses of the term heat, both now and historically

1. Support. There are so many uses of the word "heat" in technical contexts that it would be PoV to concentrate on a single one. Are we supposed to mean heat as TdS or as the transfer of energy due to a tempetaure gradient? The two definitions are not equivalent, but both are widely used. The confusion can only (IMHO) be resolved by reference to both historical and contemporary examples. Physchim62 (talk) 12:11, 3 July 2007 (UTC)

Comment: Obviously we should represent all points of view. It is against Wikipedia policy to even have a poll like this. According to Wikipedia:Neutral point of view: All Wikipedia articles and other encyclopedic content must be written from a neutral point of view (NPOV), representing fairly and without bias all significant views (that have been published by reliable sources). According to Jimmy Wales, NPOV is "absolute and non-negotiable." In other words, the history of a term is obviously a significant point of view and hence having a vote to eliminate "history of scientific" concepts (significant points of view) from Wikipedia is non-negotable. --Sadi Carnot 23:07, 5 July 2007 (UTC)

Sadi, you really do protest too much. We need to be clear, as editors, what this article is to be about. Is it to be about the concept of TdS, or the word heat? Either way, we can then write an NPOV article. But we need to settle what the article is supposed principally to be about, by properly coming to a demonstrable view, or otherwise editors will continue to scrap over it. Jheald 23:30, 5 July 2007 (UTC)

Jheald, the quantity TdS is an approximation of what heat is. This is the thermodynamic point of view, as developed by Clausius. There’s other ways of discussion heat, e.g. radiation or convection, that don’t need recourse to entropy. --Sadi Carnot 00:36, 6 July 2007 (UTC)

By TdS I was only meaning a shorthand for "heat in the narrow first-law sense". You're right that heat in that sense doesn't need to be discussed as TdS. However, the heat will in most cases still be a quantity of energy TdS, whether it is discussed in those terms or not. Hence my use of TdS as shorthand. (But not a perfect shorthand, I accept, because heat may be transferred from systems that have no well-defined temperature) -- Jheald 07:53, 6 July 2007 (UTC)

Merged heat transfer mechanisms stuff to Heat transfer article

In the years since this discussion took place, this article and the heat transfer article both came to have significant overlap and redundant information. Since it is not clear to me that they definitely ought to be the same article, I have moved the redundant elements of the heat article into the other one.

This has the effect of making this article a bit more general, and heat transfer perhaps more focused on topics specific to engineering (which may reflect my own bias). I think that it is important to note that although according to the definition of "heat" begin roughly "energy transfer" the term "heat transfer" seems redundant, nevertheless it is widely used in contexts where the word "heat" would not be equally applicable in the same sense (ie, engineering). IMO that usage is important and the article(s) should mention the overlap and potential confusion in the terminology, not try to fix it.

Note also that much of the merged material even in the heat transfer article is probably too much detail and ought to be moved to the individual articles on conduction, convection, etc. as heat transfer is actually pretty long. So the end result will simply be to focus and pare down otherwise length articles into more readable and organized pieces.

I hope that explanation makes sense...

Dhollm (talk) 12:25, 13 August 2010 (UTC)

spelling

this page says " This is a term used to characterise the combined effects of conduction and fluid flow." Now I am bad at spelling, but my spell check says it should use characterize... I don't know if it is one of those British spellings, so I will leave it be...

Robbob1508 (talk) 23:52, 13 January 2010 (UTC)

Colour?

Is there or should there be a section on the different colours when there is enough heat like white, blue, etc? I couldn't find it on any other articles. --Mooseman33 (talk) 06:40, 12 February 2010 (UTC)

Why nothing on the wavelength of heat radiation? is it "Shortwave" or Long wave radiation?

Why nothing on the wavelength of heat radiation? is it "Shortwave" or Long wave radiation? Nunamiut (talk) 14:29, 19 August 2010 (UTC)

Does not nuclear phycicists dealing with nuclear explosions talk of thermal radiation? Nunamiut (talk) 03:11, 4 December 2010 (UTC)

Temperature measures enthalpy?

In this version of the article the introduction says "Temperature is used as a measure of the internal energy or enthalpy, that is the level of elementary motion giving rise to heat transfer." As I read it, this would imply that enthalpy is internal energy, that both are entirely due to molecular motion, and that both can be entirely determined from the temperature. However,

• internal energy is just one component of enthalpy.
• molecular motion is just one component of internal energy.
• temperature measures just one component of molecular motion

The sentence therefore seems highly misleading at best, so I have removed it. (For historical context, the confusion seems to have originated from this well-intended edit) Riick (talk) 21:01, 1 September 2010 (UTC)

Opening Definition

This is not a good opening, especially when in the Overview we rely on James Clerk Maxwell's FOUR charectistics, not just transfer. Please re-write to reflect the overview and I for one, would like this to be a very readable introduction that any student looking to find about about the subject would find extremely helpful. Rjstott (talk) 16:16, 2 September 2010 (UTC)

I agree. Reading further in the overview it says "Heat flows between systems that are not in thermal equilibrium" whereas the introduction says "heat is the process of energy transfer". These two statements are not compatible, the first has heat as a property of the material proportional to its temperature, the second refers simply to a 'process of energy transfer' which seems to eliminate heat as a distinct form of energy.--Damorbel (talk) 08:23, 3 September 2010 (UTC)

Equally the first sentence in the overview "In modern terms, heat is defined as energy in transit" says nothing useful at all because it includes all forms of kinetic and electromagnetic energy, not making any distinction about the specific nature of heat. --Damorbel (talk) 09:03, 3 September 2010 (UTC)

Whatever revisions are made, I think it is important to retain at least some mention of the definition currently in the introduction - that heat is sometimes defined as the process of energy transfer due to temperature differences, rather than as the energy transferred through that process. Both definitions are considered correct, and it would be confusing indeed if a student being taught under the process definition were to find absolutely zero mention of it in Wikipedia.

Ironically, even before this discussion started I've been developing a small section regarding the inconsistent usage of the word "heat" within the scientific community. It's a fascinating topic, and definitely an important part of what "heat" is. I'll see if I can add that section today, as it may have some bearing on this conversation. Riick (talk) 14:51, 3 September 2010 (UTC)

I think I wrote the definition in the intro, and I just modified the sentence. I was not so concerned about the process of energy transfer or or the energy transfered in the process. What I was more concerned about at the time was to give the general definition, which takes into account the freedom you in principle have in defining the thermodynamic description of some given physical system.

Ultimately what we're doing when describing a system using thermodynamics is describing it statistically. The system can be a gas consisting of 10^23 molecules and we want to describe this using only a few variables. We then have to make a decision about what variables to keep in the thermodynamic description and what variables are going to be treated statistically. This choice then affects the definition of heat. So, to a certain extent, the speration between heat and work when a given amount of energy is exchanged, is arbitrary.

Example. Consider a gas in an insulated cylinder. There is a piston in the cylinder dividing the volume in two parts. We can then either chose to include the position of the piston in the thermodynamic description or chose not to do that. In both cases we're describing the same physical system, but the thermodynamic description will obviously be different. If we take the position of the piston as an additional external parameter in addition to the total volume of the cylinder, then moving the piston amounts to adding energy in the form of work to the sytstem. If we don't do this then strangely as it may sound, we're adding heat to the system when we move the piston.

To see that the latter point of view is not ridiculous, consider that in thermal equilibrium, the freely movable piston will be in an equilibrium position where the pressures on both sides are the same. If you move the piston, you're doing mechanical work on the system and then the system will be in a non equilibrium state. If the piston is released from the new postion, all the energy that was put in the system will stay there, as the system is insulated. Eventually the energy will be completely dissipated inside the system and the piston will be found in the equilibrium position.

So, when you don't include the position of the piston in the thermodynamic description, you can only describe the equilibrium states where both parts of the system are in equilibrium with each other. When we add heat to a system, there always is a perturbation away from the equilibrium state which can be made arbitrarly small. In practice one always has to wait a while before a new state of thermal equilibrium is reached as you will inevitably add the energy in a non-uniform way to the sytstem which then has to be dissipated. Count Iblis (talk) 15:43, 3 September 2010 (UTC)

In shorter words, the distinction between whether energy enters a system through heat or through work depends merely on whether that energy crosses the system boundary microscopically (eg through molecular vibrations molecular interactions) or macroscopically (eg through a moving piston rod). Thus such a distinction can vary depending on how we draw the system boundary. I agree, although clearly this does not imply that how we define the word "heat" is arbitrary. At any rate, the important thing to remember is that heat is not something inside a system. I realize you are aware of this, but I want to make sure it is also abundantly clear to anyone reading this discussion! Riick (talk) 18:57, 3 September 2010 (UTC) modified by Riick (talk) 16:57, 5 September 2010 (UTC)

I have now added the section on inconsistent usage. This is how it looks at the current time. The Herrmann source is particularly well written and explains a lot of things- anyone who has a few minutes may enjoy reading it! Riick (talk) 19:07, 3 September 2010 (UTC)

I find these fine points involving positions of pistons, whether the energy is being transferred or not etc. as quite irrelevant to the definition of heat. Heat is the kinetic energy of molecules that is transmitted between them by the exchange of momentum by elastic collisions. Elastic collisions include exchange of photons. Additional types of molecular energy such as vibration or rotation play no direct part in momentum exchange between molecules and do not directly take part in the exchange (ultimately the energy in vibration and rotation modes is equalised with that in the kinetic modes). The reason why heat has to be defined this way is because temperature is (should be) also defined the same way, for the (should be obvious) reason that heat transfer is entirely delineated by temperature gradients i.e. momentum gradients. --Damorbel (talk) 13:06, 4 September 2010 (UTC)

Temperature is not the kinetic energy of molecules, it only equals it (up to some factor), in the classical limit. The fine points you don't like are crucial to understand what temperature and heat are. To give another example: If you keep track of all the degrees of freedom in a system, then the entropy is zero and it remains equal to zero. All energy exchanges then count as work. This entropy that is zero is called the fine grained entropy. It stays zero because at the fundamental level, information does not get lost. Count Iblis (talk) 02:48, 5 September 2010 (UTC)

"Temperature is not the kinetic energy of molecules". My concern above is for the definition of heat, I referred to the necessary connection of the measure of temperature being compatible with the definition of heat. Temperature and heat are related by the Boltzmann constant which currently isn't even mentioned in the article! The Boltzmann constant is the the energy (Joules) per degree of freedom per K, it is the link between energy and temperature. To include molecules in the definition is meaningless because molecules have differing numbers of degrees of freedom, some of which only take energy at particular temperatures, the mechanics of molecules is really rather subtle, complicated and irrelevant.

Entropy is not involved in the definition of heat, entropy is a measure relating energy density (specific heat capacity) to temperature. Absolute entropy $S=k_{B}\ln \Omega \!$ is seldom calculated because of difficulty in defining the number of states at each energy level that are available to be occupied, more often it is change of entropy $dS={\frac {\delta q}{T}}.$. You will notice that the Boltzmann constant (k_B) is involved in this definition of entropy but apparently not in change of entropy, this is because Boltzmann's constant, entropy and specific heat capacity are all defined the same way J/K, most calculations of entropy change consider specific heat capacity as a constant. --Damorbel (talk) 08:23, 5 September 2010 (UTC)

What is not quite clear from the above is why entropy and temperature are not the same since they have rather similar derivations. In fact the distinction is astonishingly simple. Look at this definition of entropy from the section on statistical thermodynamics [1] _{$S=-k_{B}\sum _{i}P_{i}\ln P_{i}\!$} so entropy is a sum of the Boltzmann energies (kT) of all the particles in an ensemble (having the lnPi doesn't really affect this). From this two things can be seen, if the particles all have the same temperature then the ensemble has that temperature; if the particles are in statistical (thermal) equilibrium then a temperature can be assigned to the ensemble which is the same as if the particles all had the same temperature. From this it is easy to see how the entropy is reduced when the ensemble is in disequilibrium, the energy distribution of the particles making up the ensemble no longer has a Maxwell–Boltzmann distribution. In simpler terms, when the ensemble is in equilibrium and the entropy is at a maximum, the temperature of the ensemble corresponds to that of a single particle with the same temperature. Without equilibrium, temperature of an ensemble has no precise meaning.

From this it can be seen that when the particles are molecules and have energy in degrees of freedom that are not part of the kinetic process, the entropy of an ensemble of this kind will deviate from the model with simple particles having only three degrees of freedom, entropy and temperature will no longer be related in the same simple way.

Another situation arises when the ensemble is in a gravitational field. In this case the energy of the particles (or molecules) will depend on their position in the field, the reason for this is also fairly simple, particles moving in a gravitational field, uniform or not, exchange kinetic energy with gravitational potential energy (GPE) i.e. the change in GPE is directly changes the temperature (momentum or kinetic energy) of the particles. --Damorbel (talk) 12:46, 7 September 2010 (UTC)

Well no one changed the intro so let's have a go: In physics and thermodynamics, heat is energy which may be transferred from one body or system to another due to thermal contact. This is in turn is defined as an energy transfer to the body in any other way than due to work performed on the body.

Rjstott (talk) 12:47, 8 September 2010 (UTC)

That definition would not be correct, because if the energy is not transferred, you are dealing with a part of the internal energy. But there is no way to define part of the internal energy to be "heat" as it is not a thermodynamic state variable. Count Iblis (talk) 14:42, 8 September 2010 (UTC)

I would like some clarification of "But there is no way to define part of the internal energy to be "heat"". In paragraphs above I went into some detail about the definition of temperature and its relation to energy and entropy via Boltzmann's constant. Since temperature, a state variable, is a measure of the kinetic energy of 'particles' i.e. a measure of the energy exchange between atoms and molecules by momentum transfer, is this not the 'part of internal energy' to be called heat? If not, do you have another name for the energy of particles due to momentum? --Damorbel (talk) 07:51, 9 September 2010 (UTC)

The part of internal energy associated with particle motion is sometimes called "thermal energy" (although this term sometimes has other slightly different usages as well). Thermal energy is directly linked to temperature and entropy, as you have correctly stated, but it is not uniquely linked to heat. One could increase the thermal energy simply by doing work- by stirring a fluid, rubbing a solid, etc. Even in situations which only involve heat transfer, we still can't think of thermal energy as "stored heat". That's because when you add energy through heat, some of that energy may end up as something other than thermal energy. When you melt a solid, for example, not all of the added energy ends up as motion in newly-opened-up degrees of freedom. Some of the energy instead goes into overcoming those forces of attraction which had bound the solid in the first place. This energy is therefore unavailable for use in molecular motion and thus does not count as thermal energy. In fact, it is not possible to identify any one component of the internal energy and say "aha- so that's where all the heat goes," because in most cases some of it goes somewhere else as well. Riick (talk) 19:16, 9 September 2010 (UTC)

In Our Time

The BBC programme In Our Time presented by Melvyn Bragg has an episode which may be about this subject (if not moving this note to the appropriate talk page earns cookies). You can add it to "External links" by pasting * {{In Our Time|Heat|b00fq3d4}}. Rich Farmbrough, 03:15, 16 September 2010 (UTC).

It's better to not mention temperature in the lead sentence defining heat

I.m.o., it is better to build in a bit of separation here and explain that heat flows as a result of temperature difference a bit later (e.g. immediately after the opening definition). The problem with mentioning temperature in the definition of heat, is that it makes the concept of heat conditional on a rigorous notion of thermal equilibrium where temperature can be defined precisely. This is not necessary. Indeed, it would be paradoxical, unless you consider infinitesimal amounts of heat.

In violent process, heat transfer can happen without it being possible to define a temperature, even approximately. Cases where temperature can be defined (approximataly, in a local sense) are always special cases, albeit ones that we commonly experience. Count Iblis (talk) 17:43, 24 September 2010 (UTC)

Say what? I would have thought from definitions that you could not transfer energy as "heat" (thermal energy transfer) without the objects being at some kind of definable "temperature" δE/δS (energy change per change in klnW microstates, assuring that it's distributed in a "temperature manner" so that a certain amount of energy produces a certain amount of entropy). If you don't have a well defined temperature for the source, I can't imaging how you would "recognize" the energy transfer as "heat" transfer. You'd have energy transfer, certainly, but you'd be forced to CALL it something else, not heat transfer (yes, I know all "heat" is "heat transfer"). For example, you can absorb laser light or nice monofrequency microwaves from an oven, into a glass of water, and that's not heat transfer, even though the "thermal energy" of the water increases. It's only when you shine radiation into the water that has a "heat" spectrum (black box distribution), such as from an infrared heating element, that you have "heat transfer." And then, the heat has the "black box" temperature, so there is a source temperature, and that's it. The microwave oven and the laser produce energy with no "temperature." Each one can be converted to other types of energy without paying any entropy cost. δE/δS for both approaches infinity, so there's no limit to how hot you can heat that water using monochromatic EM waves (of any frequency). If you had an actual heat source (like using sunlight to heat the water) the temperature limit for heating anything with it, is the temperature of the source. Which in this case is well-defined as 5500 C. SBHarris 04:02, 4 December 2010 (UTC)

"If you had an actual heat source (like using sunlight to heat the water) the temperature limit for heating anything with it, is the temperature of the source. Which in this case is well-defined as 5500 C." Unless you use mirrors and/or lenses to focus the light. This is easy to show because your comment about monochromatic EM waves already demonstrates the proper knowledge that each individual photon does not have a temperature in of itself.Kmarinas86 (Expert Sectioneer of Wikipedia) ^{19+9+14 + karma = 19+9+14 + talk = 86} 14:28, 6 January 2011 (UTC)

No you miss the point. Even if you use a mirror or lens, you cannot exceed 5500 C using sunlight. WHat you see mechanically is that soon you get a small image of the sun, and cannot make the brightness within that image greater than the brightness looking at the disc of the Sun. But the real reasons are deeper than that, and it's no accident that monochromatic light is focusable down sizes that are defined by the monochromicity. Thus also the temperature, which can be as hot as the source is monochromatic. SBHarris 01:52, 7 January 2011 (UTC)

"WHat you see mechanically is that soon you get a small image of the sun, and cannot make the brightness within that image greater than the brightness looking at the disc of the Sun." If what you say in this sentence is correct, then concentrated solar power is a fraud. On the contrary though, obviously the sun is brighter on a per steradian basis if you focus the light on a tiny surface, so what you say is incorrect. To be sure, luminosity and absolute magnitude are not affected, but flux density and apparent magnitude are. (Inverse square law) Also, the wavelength itself does not matter as much, for if you can get a big enough lens, even if your resolving area is the size of a house, which is much bigger than visible wavelengths, you can have solar concentration anyway. With enough prisms, you could conceivably separate out light with desired wavelengths, including wavelengths smaller than those corresponding to the peak intensity on a 5500 C blackbody radiation curve, but this bears no relation to some preconceived limit of how much kinetic energy can be gained by bulk matter in the form of heat. It does not bear any relation to it— Proof: You could take those photons, absorb their energy into electricity, and drive an apparatus producing half as many photons (i.e. light quanta) with no more than twice the frequency, or produce a third as many photons (i.e. light quanta) with no more than three times the frequency, while always keeping in mind that $E_{photon_{in}}>E_{photon_{out}}$ if work is done or if thermal energy is stored as a result of the energy acquired by the photons. If you want to get really spicy, you could take a small section of the photons in the 5500 C black body radiation curve to the right of the peak (i.e. smaller wavelengths), and with a method similar to "cut-and-paste", replicate a much lower luminosity source with a higher black body temperature. And of course, an antenna expert could still try to absorb some of the radio frequency power to generate pretty much the same thing even though radio frequencies are on the opposite side of the 5500 C peak as the higher frequencies.Kmarinas86 (Expert Sectioneer of Wikipedia) ^{19+9+14 + karma = 19+9+14 + talk = 86} 06:47, 7 January 2011 (UTC)

You're still not getting it. The article on concentrated solar power itself points out that 5300 C is the maximum temp you can get by simple optical or reflective concentration of sunlight, however you do it [2]. That is because this is the blackbody temperature of the sun. You can transduce thermal energy to make any temperature you like (use it to make electricity to run a laser) but in doing so, the fact that you have a thermal input forces you to pay an entropy cost, so that only some of the heat is available at over 5300 C, but some of it now comes out as heat at a lower temperature, and is wasted. None of this would be necessary if heat from the Sun did not come with a limited temperature. If it came as a perfectly monochromatic laser beam, we could in theory convert 100% of it to any form of energy we liked. If it came at a higher temperature than 5300 C (actually it is about 5500 K) we could use a larger fraction of it. As it is, our maximal heat conversion-efficiency even in theory is 1-(cold reservoir temp/5500 K), and this will never be 100%.

And by the way, thermal energy is NOT kinetic energy only. That's just wrong. SBHarris 23:06, 19 January 2011 (UTC)

I certainly didn't think that entropy would not increase, in fact it would, and energy would for certain by lost from the system and become unusable. I agree with the definition offered by Friedric Herrmann which suggests that our common sense notion of "heat" is really entropy, although this has little to do with most of the article for heat. Looking at that though, I do not see any reason why entropy itself could not be concentrated onto a few particles (though obviously that means a greater number of particles have lost that energy plus the inevitable entropy increase). Perhaps there is a way to do that, but I know now that would be a discussion on quantum behavior having little to do with "classical thermodynamics".

P.S. Don't you wish they would just force the lexicon to change so that these confusions can be prevented?

P.P.S. On a last note, I thank you for you recent edit to Heat. I was most certainly wrong about equating thermal energy as kinetic energy of the particles.Kmarinas86 (Expert Sectioneer of Wikipedia) ^{19+9+14 + karma = 19+9+14 + talk = 86} 00:06, 20 January 2011 (UTC)

After a long reflection, I think the above merits the need for separate articles for heat. I suggest that one or more of the following articles be created:

• Heat (classical thermodynamics)
• Heat (statistical mechanics)
• Heat (chemical thermodynamics)
• Heat (equilibrium thermodynamics)
• Heat (non-equilibrium thermodynamics)

I hope this idea has enough merit to counteract some the disruption I may have caused here as of late.Kmarinas86 (Expert Sectioneer of Wikipedia) ^{19+9+14 + karma = 19+9+14 + talk = 86} 00:18, 20 January 2011 (UTC)

Heat as kinetic energy

I have reversed this deletion [[3]] not to have such a definitionof heat is just about as absurd as you can get. Heat as molecular/atomic motion is the basis of Kinetic theory the person trying to eliminate this connection has an awful lot of modern physics to either learn or undo. Clearly the one responsible for removing the kinetic definition has no concept of Kinetic theory and is just playing games.--Damorbel (talk) 12:23, 6 January 2011 (UTC)

Never mind that heat energy transfer is due to the acceleration of molecules and atoms with respect to each of their individual inertial frames, not merely their velocity relative to (i.e. subtracted by) the velocity of some pretentious observer.Kmarinas86 (Expert Sectioneer of Wikipedia) ^{19+9+14 + karma = 19+9+14 + talk = 86} 14:07, 6 January 2011 (UTC)

Heat as entropy

http://www.girep.org/proceedings/conference2004/Friedrich_Herrmann_-_Entropy_from_the_Beginning.pdf - This is now officially my favorite article on this subject.Kmarinas86 (Expert Sectioneer of Wikipedia) ^{19+9+14 + karma = 19+9+14 + talk = 86} 00:26, 20 January 2011 (UTC)

Equivocate?

I'm not sure what the problem is, but I reverted because I'm pretty sure a paragraph like this one is not part of the solution:

However, in informal settings, the word heat often refers to the total thermal energy Q. Consequently, certain commonly used phrases with the word heat, such as "flow of heat" and "heat flow", contain a redundancy which may reinforce some to mistakenly equivocate their personal understanding of heat as some form of energy, with the more precise heat-related concepts in thermodynamics. The precise definitions for heat, which take various forms including energy, power, and surface power density, are further qualified by distinct respective terms (i.e. heat transfer, heat transfer rate, and heat flux).

Maybe User:Kmarinas86 can explain the problem to us here. Dicklyon (talk) 01:06, 20 January 2011 (UTC)

Since editing the article, I have come to several revelations, in the following (chronological) order:

• I have discovered how (hopelessly) confused the terminology is:

• Heat is sometimes synonymous with a substance.

• At first this only made sense.

• Heat is sometimes synonymous with a process.

• I edited the article as if heat can only be treated as a process, and that any other means of treating heat is incorrect.

• Some believe that heat is naturally an entropy concept, not an energy concept.
• I discovered the article by Friedrich Herrmann via a citation.

• At first, I started with a null hypothesis that this paper was wrong.
• It was convincing enough for me to change my position on what heat actually is.
• Particularly interesting was the discussion on how the terminology got as bad as it did.
• The paper's style resonated with me completely.
• I concluded that this alternative view is the only one that makes sense.

• And then I discovered: It would have profound implications if it turned out that energy was a property of entropy, rather than the other way around, yet it seems plausible to me.
• Almost no scientists care that the concept of heat is confused.

• Obviously, any attempt to change this will be seen as confused.

So in recapitulating this, I decided that I will no longer edit this article. There is no hope for a single person in changing the stubbornness of accepting skewed definitions, whether done through scientific courtesy or not, without a complete alternative to the current theoretical formulation.Kmarinas86 (Expert Sectioneer of Wikipedia) ^{19+9+14 + karma = 19+9+14 + talk = 86} 03:00, 20 January 2011 (UTC)

It's better to edit from secondary sources, rather than let yourself be jerked around by all the opinions you'll find in primary sources. Thanks for backing off. Dicklyon (talk) 04:52, 20 January 2011 (UTC)

Personally, I think that having multiple opinions and definitions for heat is totally contrary to certain principles of science. In my view, to have multiple definitions that are in effect mutually exclusive is anti-scientific and semantically illiterate.Kmarinas86 (Expert Sectioneer of Wikipedia) ^{19+9+14 + karma = 19+9+14 + talk = 86} 13:06, 20 January 2011 (UTC)

What principles of science would those be? Why can't we just acknowledge if there are differing and evolving conceptions? And isn't there one pretty broadly accepted definition that everyone uses anyway? Dicklyon (talk) 03:25, 21 January 2011 (UTC)

"What principles of science would those be?" No knowledge should be taken for granted. Fields of research should attain semantically consistent theories. When a quantity represents some measure of physical quantity, use the physical unit for that physical quantity to denote it, not one for some other physical quantity if it is not the same. "Why can't we just acknowledge if there are differing and evolving conceptions?" Different doesn't mean evolving. "And isn't there one pretty broadly accepted definition that everyone uses anyway?" If it's the thermodynamic definition, then no.Kmarinas86 (Expert Sectioneer of Wikipedia) ^{19+9+14 + karma = 19+9+14 + talk = 86} 02:17, 23 January 2011 (UTC)

Heat, Thermal Energy and Disequilibrium

In the discussion of heat it is important to recognise, and quantify, its relatedness to thermal energy. Any system at equilibrium involves constant exchanges of thermal energy within it, but no change is induced, as all exchanges cancel out to a net zero. Now consider creating a slight disequilibrium between two parts of the system. Exchanges of thermal energy continue, but there is now an average overall direction of flow of thermal energy - this is heat. Heat is therefore the natural result of thermal energy exchange as dis-equilibrium. It is the same thermal energy exchanges that were occuring in equilibrium with a net bias in one direction introduced by disequilibrium between the two systems. Additionally heat can be excahnged "uphill" via the application of work. So heat can be exchanged (though not spontaneously) from a cold to a hot body - in contradiction to bullet point 5 of the definitions of heat.

So heat is:

- The net effect of random thermal energy exchange at dis-equilibrium. - Heat is only transfered spontaneously from a hotter to a colder body. - Heat can be forced to be exchanged between any two bodies of any temperature by the application of work. Gordion101 (talk) 12:02, 20 June 2011 (UTC)

I concur with your overall conclusions. I would suggest that the your point on bullet point 5 in the definitions is absolutely correct to what I believe is the most rigorous, generalized and common (at the expert level) definition of heat. There is however a reasonable definition supported by many in the literature that also requires spontaneity in addition to energy flux. Note that more often than not these definitions neglect rather than strictly forbid non-spontaeous heat. Although I see no loss of value in defining heat in both spontaneous and non-spontaeous directions to the group that supports this spotaneous-only definition and in the modern context of expert use the spontaneous-only definition is a minority definition, it exists nonetheless. I would encourage more clarification of the distinctions between these definitions without removing such definitions. It is also important to note that the spontaneous-only definition may appear to be more common than it actually is due to a couple of reasons: 1) by far most practical problems involving heat are in fact spontaneous, 2) the explanation of heat to the layperson is difficult to begin with and including non-spontaneous heat is a source of confusion so it is often excluded only for the sake of clarity, even at the undergraduate level.--Nick Y. (talk) 13:08, 20 June 2011 (UTC)

Thanks Nick, I take on board your point about the generality of the focus on the spontaneous side of heat. However, it is worth remembering that all forms of refrigeration rely on the "uphill" version driven by work. I remember many hours struggling over problems involving Carnot engines run in reverse etc at undergraduate level, so I think this is a reasonably important point and might be worth including to prevent omission of a serious aspect of not only the theory, but also the practical applications of heat. Gordion101 (talk) 20:21, 22 June 2011 (UTC)

I tweaked the lede just to tweek the discussion here, and remind all that you don't need mechanical work to remove heat against a gradient. All you need is something that increases entropy. A concentration change will do it, as for example the everyday removal of heat from your forehead against a gradient, by means of the evaporation of the sweat of your brow. That's the only reason you don't quicky die when the temperature is > 106 F or 41 C. Evaporation still lets you transfer heat from your body to an environment which is hotter than you are.

Of course, such a step is part of any closed-cycle heat pump, but the point is that you don't NEED to close the cycle. You can just keep drinking water. The mechanical work is only necessary for certain closed-cycle processes. For other processes you never need mechanical work, but you do need to send in energy, and it needs to be free-energy. You can (for example) cool something with a Peltier junction, if you send in electicity. You can make electricity with a concentration cell. Again that's an open cycle in a sense if you get all the way out to the cell, but then so is your refrigerator, if get all the way out to the power plant, where the zero-entropy energy of electicity is made, at the cost of dumping its entropy someplace else during production. Basically, all free-energy comes from the low entropy conditions after the Big Bang, when it comes down to brass tacks. All we ever do is manipulate the entropy in small regions, or for small amounts of energy (like electric current that runs the refrigerator) at the cost of increasing the entropy someplace else. And it is that that allows a closed-cycle heat pump to operate. SBHarris 22:48, 22 June 2011 (UTC)

Yes, these are all good examples of how the spontaneous-only definition and the hotter-to-colder requirements sometimes seen is incomplete and not a particularly logical division of phenomena. I would suggest that these definitions be reduced in prominence in the article, yet still present and well explained. It seems to me that these spontaneous-only definitions are both older historically and tend to appear more commonly in elementary contexts where intuitiveness is valued over rigor and should be characterized as such. Note that heat is actually a surprisingly challenging concept for the layperson. However, there is a minority that takes these definitions seriously and there are many expert quotes that would appear to as well (although I am certain that in many cases this is simply a lack of attention to details).--Nick Y. (talk) 13:31, 23 June 2011 (UTC)

Hi SB, Thanks for your comments. I agree that work does not need to be mechanical work. It can be any form of work - naturally including electrical. With regards to the example you give of evaporation, I don't think this is a good example of what you are maybe trying to say. In the case you discuss the sweat sits in thermal contact with the person, and then the sweat undergoes a phase change - evaporation - in which it transfers thermal energy to latent heat within itself, and hence lowers its own temperature. The body's cooling is then heat flowing spontaneously from a hotter to a colder body. There is never any heat transfered up a thermal gradient. The evaporation is, of course, an entropy driven process, that ceases when entropy can no longer be lowered by transfering water molecules from the liquid to the gas (saturation). Gordion101 (talk) 06:58, 25 June 2011 (UTC)

In the same sense we use when we speak of "heat pumps" the "heat" is transferred against a thermal gradient any time the heat is absorbed by phase change and the substance that holds the "heat" is physically moved into a higher temperature region (or moves itself, as happens in diffusion). If you condense it again after moving (say by doing work on it) the heat comes back out and you say (very loosely) that you've "moved" the heat (as in a refrigerator). But if you insist on that language, the "heat" is still there, having been moved against a thermal gradient, even if you don't ever recover it as something that raises the temperature.

Okay, perhaps you want to call it "stored thermal energy" in an evaporated fluid, and not (technically) "heat." But if you DO that, note that a refrigerator, or any sort of heat pump doesn't pump "heat," either. Istead, if you must, a heat pump actually moves/pumps thermal energy from (physically) here (inside the box) to (physically) there (outside, where temperature is higher). It's not really the "same" heat that disappears inside the box as appears outside. In one sense it's not heat at all during the moving, since by definition heat cannot flow that way, even WITH work. All that EVER happens is that heat disappears HERE and is made to reappear THERE (in a hotter place) and we pretend it's the same "heat." It never is (in your strict sense) but (again) that's true even with refrigerators. However, it is the "same" thermal energy (in the sense that energy is being conserved here), and you don't have to have a closed cycle to move this thermal energy the "wrong way" against a thermal gradient. It simply needs to be carried in some other fashion than heat flow. That's always true, work or no. SBHarris 02:28, 26 June 2011 (UTC)

I think one clear example of heat being driven up a thermal gradient would be a Peltier Cooler. Here electrical work results directly in heat being transfered from a cold to a hot resevoir. No phase changes, no mass transport.Gordion101 (talk) 20:32, 26 June 2011 (UTC)

But no true heat transfer, either. Instead you get an electric field that drives a separation of charges. Some of those charges (the hot charge carriers) go to a physical reagion where the temperature is higher, but they are driven by their charge, just like an expanding gas carrying "heat" in the wrong "direction" from a region when it moves away from it. The escape of the hot carriers removes energy from the material, and this energy comes from ambient heat, so it gets colder when the carriers have been pulled out. It's very like a phase change, followed by pulling out the evaporate. On the other side of the junction a similar energetic process causes heat to be released when the carriers appear, as in a resistor. More heat is made in the thermoelectric effect than disappeared on the other side, when "heat" is moved. However, it's not the same heat that disappeared on the other side, since actually charge carriers have been moved, just like a moving an expanded gas which "carries" thermal energy, but the physical motion of which can be called "heat" only if you expand your definitions a lot to include thermal energy carried by any means. All the same problem. SBHarris 23:37, 26 June 2011 (UTC)

The Sun and Heat

The caption for the first image on this page of the sun is not correct: "Heat generated from the nuclear fusion in the Sun, and transported to earth as electromagnetic radiation..." Heat can never be "transported as electromagnetic radiation" - which is a form of work. The opening definition of heat clearly expresses the fact that heat is energy transfered through thermal contact and specifically excludes energy transfer via work. The sun has 150 million km of vacuum between it and the earth - no heat is going to get through that. Additionally the definition of heat is about transfer, so the first part of the sentence - that heat is generated in the sun - is also wrong. Thermal energy is generated through conversion of nuclear potential energy to radiation (which is subsequently absorbed) and kinetic energy in the centre of the sun, but this is not heat. Heat is thermal energy transfered by non-work means to another system. Basically the whole caption is completely in contradiction to the rest of the article.Gordion101 (talk) 19:58, 25 June 2011 (UTC)

I think most people would count black body radiation as "heat" and radiant transfer of heat as, well, radiant heat transfer (it's a whole engineering subject). Nor do I think that EM radiation is normally counted as "work." It is a form of energy, of course. But work is force times distance, since by "work" we ordinarily mean "mechanical work." If you look at work (physics) mechanical work is all you find. They are synonymous. SBHarris 00:51, 27 June 2011 (UTC)

Heat, temperature and energy

So long as the opening line is 'In physics and thermodynamics, heat is energy transferred from one place in a body or thermodynamic system', the article is always going to be misleading. The measure of heat is temperature, i.e. energy per degree of freedom, it is measured in Kelvins, degrees Fahrenheit, degrees Celsius etc; energy (by itself) is measured in Joules, ergs, calories etc. The important distinction can be seen in that two equal volumes (or two equal masses) with the same temperature may possess different amounts of energy. As currently written the article makes nonsense of the 2nd law of thermodynamics; personally my money is on the 2nd law --Damorbel (talk) 06:42, 26 June 2011 (UTC)

Heat is the form of energy that flows between two samples of matter due to their difference in temperature. Temperature, in turn, is a sort of measure of energy concentation into a small number of states, and specifically the concentration of internal energy per unit of entropy in a system. So yes, I see no way of talking about heat Q without starting to talk about entropy right away. Entropy is tied up with disorder: perfect order gives an entropy of zero. Entropy also tells how much of a certain fraction of a quantity of energy can be used to perform mechanical work (PV or F*d), since entropy (disorder) must always increase in a spontaneous process. If entropy is zero, then all the thermal energy can be turned into energy and (almost) none need be left to maintain the entropy. If, however, there is some entropy S in the system, then some of the energy T*S is trapped as thermal energy, unless some way or other can be found of paying its ransom, by leaving the system with an equal amount (or more) of disorder (such as mixing atoms or spreading them out in space).

If all of a certain amount of a quantity of energy E can potentially spontaneously be turned into work (such as is true of a single photon, or a sytems of photons all having the same frequency and energy), then the entropy of the energy is zero, and its effective temperature T is dE/dS, which is infinite. Otherwise, introduction of disorder into the perfection of the way energy resides (either as photons or motions of atoms, which must all be in the same direction if they exist at all) generates entropy, and there is no fixing this. Once done, it is permanent, and the energy is degraded in a way that prevents some of it from ever being turned into a pure coherent beam of light (or any other electromagnetic energy), or completely into any type of potential energy than can be turned into mechanical work (such as, for example, electrical power or chemical potential) or into mechanical work itself.

It's amazing how difficult this is to explain, as the concepts are all interlocking. SBHarris 01:27, 27 June 2011 (UTC)

I would like further explanation of your statement 'Heat is the form of energy that flows between two samples of matter due to their difference in temperature'. I do not understand how you can say 'Heat is the form of energy' when heat, i.e. temperature, is not measured in Joules (calories ergs etc.) but in K, ^oF or ^oC. Further problems arise when introducing entropy ('Entropy is tied up with disorder') which is an extensive (bulk; macroscopic) property as compared with heat which is an intensive (local; microscopic) property.

Let us remind ourselves, the key definition of temperature is the Boltzmann constant k_B. If you check this link http://arxiv.org/ftp/arxiv/papers/1102/1102.4831.pdfyou will notice that the Boltzmann constant is most likely to be used in the near future as the fundamental physical defintion relating temperature and energy, replacing the Kelvin. By contrast, entropy is a property which reaches a maximum only when the material concerned has a uniform temperature, entropy can only be lowered when there is a temperature difference between different regions of a thermal system - makes it hard to use 'entropy' for defining temperature. Looked at another way, temperature is a property at a microscopic level - that level being the 'degree of freedom' (DOF). If a thermal system has an entropy below the maximum then it means the system has different temperatures at different locations. A thermal system only has a single temperature if that temperature exists throughout the system i.e. it is in equilibrium and its entropy is at a maximum. --Damorbel (talk) 18:05, 12 July 2011 (UTC)

Excuse me for interrupting this high-brow discussion with a low-brow concern: In the final paragraph of the lede, it says “Heat is also often referred to as thermal energy”. But shouldn't that be the other way around? I.e., “Thermal energy is also often referred to as heat”? That seems to be what the rest of the sentence is all about: The common (mis)use of the word “heat” for thermal energy. Hanche (talk) 18:58, 16 August 2011 (UTC)

I do agree Hanche, if you keep it simple - 'heat = how hot' and thermal energy = 'how much hotness', the picture will stay suitably simple. But introducing expressions such as 'Heat Pumps' and 'The transfer of energy by heat from one object to another object' etc., etc. is just confusing nonsense.

I don't demand the exact words 'hotness' and 'how much hotness' but they are a lot closer to the meaning of 'heat' than all the words currently comprising the article.--Damorbel (talk) 20:04, 16 August 2011 (UTC)

No, no, no. Heat is NOT the same as temperature. Heat is thermal energy in motion. Which is why heat and thermal energy have the same units, but temperature (hotness) is something else. A glowing spark has more "hotness," (high temperature) but a bathtub of water has more thermal energy. To heat the bathtub of water takes the transfer of a great deal more heat. Due to its far larger heat capacity. SBHarris 22:23, 16 August 2011 (UTC)

'Heat is NOT the same as temperature' What then is temperature a measure of? You write 'Heat is thermal energy [in motion]', but energy is measured in Joules. Temperature; heat; hotness etc. in measured in Kelvins, ^oF; ^oC etc. not in Joules.

Energy in motion (your heat?) is measured in Joules per second or Watts. These concepts and definitions are at the heart of thermal physics, it is not pedantic to insist on them; put correctly in the article they will free it from a lot of confusion. --Damorbel (talk) 08:46, 17 August 2011 (UTC)

I'm not sure what your beef is, at this point. By convention, we require heat to be thermal energy transferred across a boundary (otherwise, if it is just residing in an object and not moving, we call it "thermal energy," not heat). So heat that moves into an object becomes thermal energy once it is there. However, although heat is required to be thermal energy transfered along a temperature gradient, heat is not defined in units of power. Rather, heat Q is given in energy units (joules) and heat transfer rates (dQ/dt) are in units of power (watts). Neither joules nor watts as SI units are capitalized (though J and W are capitalized).

Finally, temperature does not have these units, since it is measured in Kelvins and so on, but this amounts in the end to an average kinetic energy per particle for a system in thermodynamic equilibrium. For radiation fields and systems not containing massive particles, the temperature can be defined by extention (if it exists) by comparing it to systems of particles that do have a temperature. SBHarris 16:54, 13 September 2011 (UTC)

You write "otherwise, if it is just residing in an object and not moving, we call it "thermal energy," not heat" My problem (with the article) is that heat is measured by temperature e.g. ^oC; ^oF or K and thermal energy, like all forms of energy, is measured in Joules (J).

I am not sure what you are referring to when you mention a 'convention' requiring 'heat to be thermal energy transferred across a boundary'. Is heat not a property 'just residing' in all material above absolute zero (0K), on any side of a boundary? --Damorbel (talk) 17:29, 13 September 2011 (UTC)

Heat is NOT measured by "temperature." Only temperature is measured by temperature! If you stick a thermometer in something, it does not tell you anything about how much "heat" is in it. The difference between "heat" and "thermal energy" is more subtle. Not very long ago, the two terms were used interchangably, so that we could reasonably talk about how much heat there was "in" a bathtub of water (answer: cool it to absolute zero by contact with a cold reservoir, and measure how much heat came out). Nowadays, thermodynamicists are trying to reserve the term "heat" for thermal energy that is transferred. Thus, it would be correct to say that you can find out how much thermal energy there is in a bathtub, by finding out how much heat you need to extract, to decrease its temperature to absolute zero. However, a bathtube twice as large will have the same temperature, but twice as much thermal energy, and you'll need to extract twice as much heat. Do you see? SBHarris 17:39, 13 September 2011 (UTC)

"Only temperature is measured by temperature!" In non-thermodynamic parlance a distinction is made between 'sensible' heat (that which you can feel) and 'latent' heat. Sensible heat is characterised by a temperature change whereas latent heat does not generally involve a temperature change; melting ice and boiling water being the commonplace examples. What is measured by temperature is that energy associated with the kinetics of particles. The whole point of latent heat is that it is not associated with the kinetics of particles: for ice to melt there must be energy supplied to overcome the forces holding the molecules rigidly in their (crystaline) positions: to convert water to steam the attractive forces between water molecules (holding them in groups) must be overcome; as you add thermal energy to a volume of water the ratio between liquid and gaseous state gradually shifts from - 100% water (boiling point has just been reached) to 0% water (finally the temperature begins to rise again - the steam is now 'dry' and becomes 'superheated' as the temperature of the molecules starts to rise). The important point being that the temperature (of the molecules) does not rise while the water is changing state. So here you have a vast change in molecular energy (ice to water - water to steam) with no change in temperature. Surely there should be some way of distinguishing this energy from heat? Not making a distinction seems very unscientific! Without a distinguishing measure of temperature the second law of thermodynamics will look rather stupid. --Damorbel (talk) 20:51, 13 September 2011 (UTC)

Since a lot of any ordinary "heat" (or in modern terms, thermal energy) is composed of potential energy (in every case but ideal monatomic gases), there's no point in coming up with a different phrase for the "latent heat" that disappears into phase changes without a change in temperature. Or rather, we already have: it's called "latent heat" as you say yourself. It goes entirely into potential energy (not kinetic energy) and that's why it doesn't affect temperature. What's the problem? SBHarris 21:24, 13 September 2011 (UTC)

The problem arises from calling potential energy latent 'heat', there is nothing 'kinetic' about the energy of state change (fusion and evaporation) that is why the temperature does not change and it shouldn't be called heat. The 2nd law is all about changes in (kinetic) energy arising from differences in temperature. It is true that energy changes like fusion arise from temperature differences but so do other kinds of energy change like chemical reactions, sometimes confusingly called 'heat of reaction' when energy of reaction would be more appropriate. 'Temperature' and 'heat' are all about the kinetic energy of atoms and molecules and must be distinguished from other forms of energy. --Damorbel (talk) 05:58, 14 September 2011 (UTC)

I think we've come to the source of the problem: you're wrong about the definition of heat! Heat energy (now called thermal energy) is usually not just kinetic energy; the only time it actually is only kinetic energy, is in thermal energy of ideal monatomic gases (as I said above). Heat that goes into a phase-change (heat of vaporization, etc) is most definitely a type of heat, even though it does not change temperature. You noted that it is sometimes called latent heat. You have opined that it shouldn't be called heat, but it universally is called heat, so the community of scientists disagrees with you. When the vapor deposits, and turns to liquid, it transfers heat to the object it deposits upon. Half of the "heat" (okay, "thermal energy") stored in solid objects as a result of their temperature (at least those that reach the Dulong-Petit limit of heat capacity) is due to potential energy of atoms in these objects, and that is why such solids have twice the the atom-molar heat capacity of monatomic gases-- they have another way to store heat that is NOT due to motion of atoms (rather, it resides in the fields that confine the vibrations, rather like energy in a pendulum, or a spacecraft in an very eccentric orbit).

Finally, you are completely wrong about the second law being about changes in kinetic energy arrising from differences in temperature. The second law is about changes in availablity of quantum states in phase-space. Sometimes these have nothing to do with kinetic energy or temperature. A simple example is a gas expanding into a volume, without doing work: its entropy increases, but its temperature and the kinetic energy of its molecules do not change at all. If your world view of entropy does not encompass such entropy changes that have nothing to do with heat, temperature, or kinetic energy, then your view needs modification. Wikipedia is here to help educate you, but we assume that you've made an effort in other ways first, since this is not Wikiversity. Finally, read the WP article on heat capacity. It will help. SBHarris 18:41, 14 September 2011 (UTC)

This leaves you with an interesting question, why is the Boltzmann constant expected to replace the Kelvin as a fundamental constant? The Boltzmann constant relates the a measure of kinetic energy in a degrees of freedom to temperature (K) thus it does not include potential energy of any sort.

I just do not understand how you can relate energy exchange without temperature change to heat. What about batteries? Charging a battery fills it with energy but this does not raise its temperature. Has it been 'heated' according to your definition of heat? --Damorbel (talk) 21:15, 14 September 2011 (UTC)

The answer to the last question is of course no, since the energy that goes into the battery does not go into the various expansions of freedom in phase space that cause an increase in entropy-- with the exception of the entropy of the chemical reaction itself. So there often is a small entropy cost, but hard to say what direction it is in-- this depends on the type of battery/cell. Which is why you can (given the right chemical reaction) extract the energy of a charged electrochemical cell as free energy with 100% efficiency (in theory-- in practice it's still far closer to perfect than any heat engine). However, if you use the energy that you use to charge a battery to (say) vaporize some water, some of that energy is now degraded, due to the TΔS entropy cost of vaporizing the gas. Now you can't get it all back-- even if you use the steam to drive a perfect turbine, you only get a fraction of your input energy back. So your energy goes to heat in the first case (you heated the water to vapor/steam), but has not been turned into heat in the second (where you used the energy to charge a battery which did not degrade the energy as it happened). Remember that ΔG = ΔH - TΔS. ΔG is related to the voltage and the energy you can get back out of your battery, and ΔH is the enthalpy of the chemical reaction or the phase change. If ΔS is zero or negative, as in an electrochemical cell, you're in good shape and ΔG is about ΔH. However, in a phase change like a vaporization ΔH is TΔS (they are equal at equilibrium) and your ΔG is thus going to be low, even under non equilibrium conditions (when you run your turbine you are going to drop T to condense the steam, so you can work with that). Electrochemical cells store energy as a sort of potential (electrical potential), but it's not the kind of potential that spreads out energy in all "directions." In a sense, it only goes in ONE direction, and is thus recoverable. Not so with the bits of energy that spread themselves out into the various potential states in a heated object. These are random in direction and amount, and that randomness has a price-- it can't be fixed. This energy has now been degraded to the equivalent of "heat" at this temperature. But that degradation hasn't happened yet, when you charge a cell.

Which brings us to your second question, to which I simply observe that degrees of freedom apply equally to potential energy degress of freedom, when you're talking about vibrations in a lattice in a heated object. In a gas, each degree of spacial translation (kinetic energy) represents 1 degree of freedom, and so the heat capacity of a monatomic gas is 3/2 * kT per atom or 3/2 * RT per mole of atoms. But in a Dulong-Petit solid, now each direction in space gets 1 degree of kinetic energy freedom and 1 additional degree of potential energy freedom. So now, it takes twice the heat to get to the same temperature in a mole of atoms of solid: 3R not 3/2 * R. The extra heat is stored in potential energy degrees of freedom. You haven't read the heat capacity article yet, have you?

I should add something at this point, and ask you to remember that not all kinetic energy is "heat"-- only degraded kinetic energy. A beam of gas atoms all moving in exactly the same direction and velocity has an entropy of zero-- it might as well be a perfect crystal, except for the reference frame (and you could find one where it isn't moving at all) and as such it has no temperature and no heat content (thermal energy) no matter how much kinetic energy it has, in the reference frame you choose. All that (the kinetic energy) is completely relative! But aim it at a hole in a box and once it's inside, close the door behind it and allow it to bounce around until it has attained complete randomness of direction of molecules, and a Boltzmann distribution of kinetic energies (pretend it's an ideal gas) and now it WILL have a thermal energy content. And that will be its former kinetic energy, which now cannot be made to go away by frame choice, and is a part of the system's invariant mass). And it WILL have a temperature in your new box (given by 1/2 * MV^2 = 3/2* RT, where M is the mass of the gas, V its beam velocity with respect to the box, and T is the new temperature that didn't exist before). In order for kinetic energy to be thermalized, it must first be "randomized." Do you see? SBHarris 21:58, 14 September 2011 (UTC)

"With the exception of the entropy of the chemical reaction itself." You cannot separate entropy into different kinds. Entropy is about the distribution of energy of all kinds, not just "thermally related" energy. This should be evident from the property of physical systems to change the nature of the energy, of any sort, in the course of any process, be it freezing, boiling, chemical or what have you. The importance of temperature is that, together with the concept of heat, it is the (thermal) parameter that decides which which direction an energy exchange will go. In this respect temperature corresponds with voltage which does much the same for electrical energy.

"You haven't read the heat capacity article yet, have you?" I have. But even that article doesn't mention yet another, and very important degree of freedom, the gravitational potential in the atmosphere, responsible for the lapse rate.

It is interesting to consider the kinetic energy of "A beam of gas atoms all moving in exactly the same direction and velocity has an entropy of zero" Since you have defined the entropy as being zero it is hardly surprising when it turns out to be the case. A much more interesting case is the antitank solid shot this missile converts the kinetic energy arising from a chemical reaction into kinetic energy "all in the same direction" in the barrel of the gun. When this projectile, not necessarily at a high temperature itself while travelling, reaches the target, then this "KE of flight" is rapidly converted into "KE of heat", causing chemical changes in the tank crew.

From this you should be able to see that temperature is only associated with the microscopic kinetic energy called heat and represents the intensity of the energy in a particular place which is the kinetic degree of freedom.--Damorbel (talk) 06:30, 15 September 2011 (UTC)

Bad example on your part. An antitank round does NOT convert but a tiny part of chemical energy of the propellant into kinetic energy. An artillery gun or any such device is a simple heat engine no different from the piston in the cylinder in your car, and works the same way. In your car a small fraction of heat of burning (20%) is converted to motion, and in an antitank weapon, less. That's due to the inefficiency of getting work out of heat, and THAT in turn is due to the second law. Basically, a lot of chem energy (heat) remains in the expanded gas. OTOH, if you have a fuel cell running the car, or a railgun projecting the shell, you can do a lot better, and that's because the energy is not being extracted from heat, so you don't have to worry about entropy. Electrical energy (electrical potential from an electric field) can be converted perfectly into motion, and vice versa. The same for a gravitational field. A book falling from a table converts grav potential to kinetic energy at 100% efficiency. and if an object in motion moves upward, the same happens in reverse. But if you allow any part of this to go into heat, now all bets are off, and some must stay in heat (or other forms of degraded energy) forever. No more conversion to other types of energy is possible. SBHarris 16:57, 15 September 2011 (UTC)

Time out

Damorbel, can you concisely recap what are your concerns regarding the article? VQuakr (talk) 06:39, 15 September 2011 (UTC)

The following facts are not mentioned in the article:-

Heat is measured by temperature (T).

Heat is the kinetic energy of a (fixed) number of (microscopic) particles. Heat is not energy in general, it is the kinetic energy only.

Temperature (T) is that kinetic (heat) energy divided by the number of particles in random motion (related by the Boltzmann constant), temperature is therefore a measure of the kinetic energy density; it is thus independent of the number of particles.

Heat, like temperature, is an intensive property.

"Energy" is proportional to the number of particles, it is therefore not a 'density' but an extensive property.

"Energy" is, by definition, not confined to kinetic energy, "Energy" includes potential energy, chemical energy, electrical energy etc. thus these forms of energy must be excluded from the definition of heat. --Damorbel (talk) 08:30, 15 September 2011 (UTC)

The article quite rightly doesn't, and shouldn't mention any of these things, because they are all incorrect.

"Heat" is not just a measure of how hot things are, as you will find if you check in thermodynamics textbooks. (And remember, in the event of any dispute, what we write here should be sourced to authoritative sources -- so if you disagree, you need to provide sources and to quote from them to highlight exactly what it is that they say that you think supports your assertions).

"Heat", Q, is defined as energy that is being transferred to or from a system in microscopic properties of the system, that is not reflected in the work done by or on the system that can be calculated from the macroscopic variables like pressure and volume.

"Heat" is not the average kinetic energy of the molecules of the system. If you have been taught this, then your teaching was simplified and incorrect, and if you persist with physics you will find you will subsequently be taught something different. A quantity in some ways related to what you're thinking about is internal energy -- but read that article to see that it's a rather different, more inclusive quantity. In general, it doesn't make sense to split up internal energy into a macroscopic (available) part and a microscopic (random, unavailable) part, because the availability of that energy to do work depends on how cold a reservoir you can dump entropy into. If you have a cold enough reservoir, you can extract almost all of the internal thermal energy of the system as work, with only a little bit being "wasted" as heat; but if you are working with only a slight temperature difference, then you can't do very much work at all unless you also transfer a lot of energy as heat -- where here we are using the word "heat" in its proper sense, of energy being transferred to or from the system.

Only under special circumstances can you define temperature as proportional to the average kinetic energy of particles. When this works, it is a consequence of the equipartion theorem of energy. But the equivalence only holds for independent quadratic degrees of freedom -- which might be true for an idealised ideal gas, but starts to fail to be true for real gases (and even more so for liquids and solids) as the temperature is reduced, interactions between the particles increase (so they are no longer independent), and quantum effects have to be taken into account, which can start to partially "freeze out" particular degrees of freedom. So again, what you have been taught about temperature is simplified and introductory -- something which is almost true under certain circumstances, but not generally true. Indeed, there are some systems that have temperatures that are very differently related to their kinetic energy -- for example systems of electron spins. A more fundamental definition of Temperature is that (1/T) = (dS / dE) -- ie Temperature is inversely proportional to the rate at which the entropy of a system increases as you add energy to it.

Finally, heat is not an intensive property of the system, because it is not a property of the system. Heat, as technically defined, is a quantity of energy being transferred. It is proportional to the amount of energy being transferred, so is an extensive quantity.

I am aware that this may not correspond to how the word "heat" is used loosely in everyday speech -- the loose everyday sense in which it may have been used even in high-school science classes -- simply to mean how hot something is. But "heat" in physics and engineering, particularly in the context of things like the first law of thermodynamics, has a specialist technical meaning; and it is that specialist technical meaning -- of an amount of energy being transferred -- that this article is about.

Hope this helps, Jheald (talk) 10:05, 15 September 2011 (UTC)

Heat is not measured by temperature? I say it is, what is your definition of temperature? My definition of temperature is based on the Boltzmann constant that relates temperature to microscopic energy, you can read about it in the "Report to the CIPM"here which explains why the Boltzmann constant should be selected as the basic unit relating energy to temperature. You write "Only under special circumstances can you define temperature as proportional to the average kinetic energy of particles." I don't find any special circumstances in the CIPM Report, would you care to comment? If you don't accept the CIPM position on the relationship of temperature, energy and the Boltzmann constant, then further discussion is rather a waste of time, don't you think? --Damorbel (talk) 11:37, 15 September 2011 (UTC)

If you look at the CIPM report, you will find that the definition of temperature that I gave above occurs as equation (3) in the report. It is the standard definition of temperature in classical thermodynamics, and one which the paper takes as a starting point in proceeding to define units.

Now statistical mechanics gives a precise theoretical formula for the entropy: S = k ln W; or, more generally, S = -k Σ p ln p.

In defining units, taking equation (3) together with this formula, and building on the fact that a accepted scale of units is already in place for energy, one therefore has a choice: one can either define a temperature scale (for example, in terms of the triple point of water), and then regard the numerical value of k using this scale as a measured quantity. Or, one can regard the value chosen for k as what is fundamental, making the numerical value of the triple point of water the observed quantity.

CIPM have recommended the second. This recommendation is similar to defining the numerical value of the speed of light as a defined constant, so that the definition of the metre is then determined by the definition of the second; rather than defining the metre in terms of the length of a particular bar of platinum and regarding the numerical value of the speed of light as an observed quantity.

As for the "special circumstances" required for a definition of the unit based of average kinetic energy, these occur explicitly in the CIPM's candidate definition 2: "atoms in an ideal gas at equilibrium" -- not general particles in general materials. CIPM note that one could try to work around this point, by presenting candidate definition 3 based on the context of the equipartition theorem -- but they note that this has its own difficulties, as well as requiring the concept of an accessible degree of freedom. So ultimately the report deprecates such a definition, preferring the arguably more fundamental definition 4, that the temperature scale is fixed so that the Boltzmann constant should have a particular value.

Two things about the report are perhaps particularly worth noting:

(1) It is about the definition of a unit of temperature, rather than the concept of temperature. In fact the definition of the unit proceeds on the basis that the general definition of the concept is well understood, so one can (if one chooses, though ultimately they don't recommend it) use language that fixes the definition of the unit based on a particular special circumstance, because the underlying concept of temperature is sufficently understood that one could relate other circumstances to a particular special circumstance used to define the unit, and be able to say whether the two set-ups were at the same temperature.

(2) The only places that the word "heat" appears in the document are on pages 4 and 5, where it is used in the context of amounts of heat entering or leaving a system, exactly as presented above by everyone apart from you. Heat is not measured by temperature; heat is a quantity of energy, measured in Joules. Jheald (talk) 12:47, 15 September 2011 (UTC)

"heat is a quantity of energy, measured in Joules" - that is the meaning of the article's opening statement when it states "Heat is energy". So heat is not measured in K, ^oF, ^oC etc. it is "measured in Joules". So when you go to the doctor he says 'your temperature is x Joules' does he? What you are missing can be seen in your interpretation of equation (3) in the report. Your intepretation of 1/T = dS/dQ - let us get temperature directlty by writing it as T = dQ/dS; this is the ratio of two terms dQ and dS, both of which are related to energy (Joules), so if T is a ratio of two energy terms; how can, in your interpretation, T itself, be measured in Joules? Temperature is a ratio between the actual energy of one fundamental unit, which is the kinetic degree of freedom of a particle (DOF)(particles have 3 or more kinetic DOF). The Boltzmann constant, soon to be a fundamental constant, relates this energy to any scale of temperature we care to use, like I said K, ^oF, ^oC

"S" is not an energy term. It doesn't have units of energy, but of heat capacity. Your argument fails right there, in the basic math of the units. If you want to live by an argument, you'd better be prepared to die by it, and you're wrong here. SBHarris 15:22, 15 September 2011 (UTC)

This is my original objection to the article, it stands up very well. --Damorbel (talk) 13:50, 15 September 2011 (UTC)

If you want to discuss temperature, go to Talk:Temperature. As I wrote above, in physics and engineering the word "heat" has a specialist technical meaning. This article is about heat in the sense of that specialist meaning. In physics, heat, Q, is not Temperature T. In particular heat, Q, is not a measure of how hot something is. This article is about "heat" as the word is used in physics -- for example, as used when people discuss the first law of thermodynamics. If you want to discuss some other concept, you're in the wrong place. Until you are prepared to understand that, there is no point in further discussion with you here. Jheald (talk) 14:14, 15 September 2011 (UTC)

"If you want to discuss temperature, go to etc" The original point I made was about the first line where it says "heat is energy transferred from one body.... &c" Heat and temperature are very closely intertwined in the concept of thermal energy, it is quite impossible to separate the two. Since the 2nd law describes how energy is transferred in thermodynamic systems such specialisation surely will lead to confusion, don't you think? --Damorbel (talk) 17:16, 15 September 2011 (UTC)

First off, thank you for the recap. Now that I understand your points though, this is not an ambiguous case - your concerns are incorrect. Heat is measured in Joules, BTU, or any other unit that can also be used to measure work, this can be verified from the first few pages of any thermodynamics textbook. Your CIPM reference is not a great reference for what you are trying to say because as others have pointed out, the article skips some important background since its focus is on the definition of the Kelvin, not on heat. Heat energy, Q, is clearly an extensive property as it is proportional to the mass involved; you might be thinking of molar heat capacity, which is an intensive property. While discussion of talk pages is important to developing an encyclopedia, knowing when to drop a dead issue is valuable as well. VQuakr (talk) 15:24, 15 September 2011 (UTC)

VQuakr you write "Heat is measured in Joules" I think not. If this were true, what then does temperature measure? Think carefully, it is temperature that drives the 2ndlaw of thermodynamics. It is the 2nd law that governs the direction of energy (measured in Joules) tranfer in a thermodynamic system, so it is a very important scientific concept.

You write "Heat energy, Q, is clearly an extensive property ..." How so? Heat is a function of temperature x C, a constant which is intensive.

You remark on the similarity of work and thermal energy which is interesting, it has its limitations both represent energy but heat is thermal energy/microscopic concept and is a function of the heat capacity of a substance and its temperature, so heat (thermal energy = Joules) = Q = thermal capacity (C = Joules/K) times T(temperature = K). Work is a macroscopic concept most commonly given as W = F x S (work = Joules: F = force times S = distance). Although both work and heat are measured in Joules, in practical terms they are quite different and one cannot be converted to the other with 100% efficiency. It has long been a a difficulty to reconcile heat and energy the stumbling blocks always being the 1st & 2nd law. --Damorbel (talk) 17:16, 15 September 2011 (UTC)

Work can converted to heat with 100% efficiency. Heat can even be converted perfectly to work, so long as the entropy "cost" (due to loss of heat from the universe) is paid in another way (as in a concentration cell). Heat energy is extensive because it depends on mass at the same temperature. There is more heat energy in a bathtub of water than a cup of water at the same temperature. Thus, heat energy (thermal energy) is extensive by definition. Temperature, as you've been told many times, measures the average kinetic energy per particle (which is what makes it intensive) of a system in thermal equilibrium, divided by a constant to change the units (k or R) plus yet another value (NOT a constant, but a free, measurable unknown parameter) that takes care of effective number of degrees of freedom in a system, which tells one how much of the total system energy manifests itself AS the type of randomly-directed kinetic energy we define as temperature. This last one allows changing k or R, to S. Since the degrees of freedom can vary widely according to the physical system, there is no direct connection between temperature and energy. You cannot simply equate them, because their relationship depends on the system, and is diffferent for each system. That is merely the same as saying that heat capacity is not a constant, but varies widely between 0 (at 0 kelvin) and a quantity which is more than 3R per mole (as various electronic, nuclear, and other degrees of freedom come into play) at higher temperatures. SBHarris 17:52, 15 September 2011 (UTC)

Additionally to what SB Harris has written, for your (Damorbel's) information please note that heat capacity, C, is an extensive quantity. It is the specific heat capacity c, ie the heat capacity per unit mass, which is an intensive quantity.

The relationship between energy and heat, as expressed for example in the efficiency η of a heat engine, has been well understood since at least the 1860s. There is no mystery here.

Are you even sure what you're trying to argue any more? In one paragraph you write "'Heat is measured in Joules' I think not." In the very next you then write "Although both work and heat are measured in Joules..." -- seemingly accepting the point you've just set out to deny.

Can I suggest if you want to take this further, you please first go away and read up on the subject from what Wikipedia would consider a reliable source. If you think there is a contradiction between what you have read here and the reliable source, then come back here. People have gone out of their way to try to set you straight here; but at the end of the day, as ARBCOM has ruled (eg here), talk pages are not the places to discuss your own original misunderstandings. Limited discussion for clarification can be harmless, even helpful; but when discussion goes on and on, and it is clear that you are not taking on board the points being made to you, then it is time to draw a line. What matters here is not what you think, but what reliable sources say. So either frame your comments (without synthesis) in terms of what some reliable source has to say, that you think is not properly being taken account of in this article, or find some other forum to air your misunderstandings. Jheald (talk) 17:59, 15 September 2011 (UTC)

Wups, yep. When I wrote heat capacity above as varying from 0 to more than 3R per mole, obviously I mean specific heat capacity (which is the "per mole" part). It's 3R per atom-mole. The naked heat capacity of an object can be as large as you like, since the object can be as large as you like. The specific heat capacity is more constrained, but still varies widely between zero and some maximal value where every particle participates in every possible degree of freedom of energy storage. SBHarris 18:08, 15 September 2011 (UTC)

Sorry I'm on holiday for a week or so, but in the meanwhile just remember to think about what is measured by temperature and why it is important in the 2nd law of thermodynamics. You could even think about why systems full of thermal energy are quite unable to produce any work unless there is a temperature difference. --Damorbel (talk) 18:18, 15 September 2011 (UTC)

That statement is incorrect. You didn't read the article on concentration cell, did you? Electricity is produced. Temperature difference can be zero, though the cell does absorb heat to turn into electricity (however the gradient can be as small as you like, and is not important). SBHarris 18:46, 15 September 2011 (UTC)

And the relevance of that comment to "heat" and the improvement of this article is suuposed to be what, exactly ? Jheald (talk) 18:27, 15 September 2011 (UTC)

"Temperature difference can be zero." Does it have to be positive? Couldn't an infrared laser could be used to heat up a surface of a UV light more so than the other way around? You could have multiple lasers imitating a black-body spectrum in the infrared. There can also be a variety of the lights (more than just UV lights) that imitate a black-body spectrum. But according to the laws of thermodynamics, wouldn't it be assumed that the UV lights heat the infrared bulbs, not the other way around? Is there a limit to the temperature that can produced in an object based on the frequency of light alone, regardless of the amplitude? For example, is it impossible to heat something above a black body temperature of 4000K using only high amplitude infrared light? What happens if the emissivity of the infrared lasers is much higher than that of the UV lights? Can we have a asymmetric relationship where the UV lights get hotter because of the infrared lights? Or are they unable to absorb infrared and therefore not allow infrared to heat it?siNkarma86—Expert Sectioneer of Wikipedia
^{86 = 19+9+14 + karma = 19+9+14 + talk} 22:35, 16 September 2011 (UTC)

Too many questions. Monochromatic sources don't have "temperatures" no matter what their wavelength. For this reason their energy is available to do anything. An infrared laser can be tripled in frequency to the UV (this is done in nuclear fusion experiments). Or a microwave oven can be used to heat a small sample of metal red hot. A monochromatic source is like our ideal beam of gas atoms all going the same direction at the same speed. Because "heat content" requires that something have attained a temperature, these sources don't really have heat contents, either. They only have energy contents. A black body radiates "heat," because the energy comes in many frequences and there is randomness in that. But a laser does not. SBHarris 02:57, 17 September 2011 (UTC)

"A monochromatic source is like our ideal beam of gas atoms all going the same direction at the same speed....A black body radiates 'heat,' because the energy comes in many frequences and there is randomness in that." This sure would imply a limitation of classical thermodynamics in its ability to describe systems which do not have a definable temperature. Also, I don't see a lot of discussions about the nature of "heat" inside atoms despite the inherent randomness in them. How could they be a heat source during chemical reactions if they were not hot? Do they somehow delay their emission of heat as a result of low emissivity at extremely high temperatures (frequencies)? How is that even supposed to work? Or is there no "black-body" inside an atom where the energy is structure is a very low-entropic sort of a way?siNkarma86—Expert Sectioneer of Wikipedia
^{86 = 19+9+14 + karma = 19+9+14 + talk} 11:33, 17 September 2011 (UTC)

You have to stop thinking of "heat" (we now call it "thermal energy" unless it's in motion between two reservoirs along a temperature gradient) as being ONLY energy. It's energy, but of a special type: THERMAL energy. That means energy which has been degraded by being spread out and filling a lot of possible quantum states in a system that can be energized. Example: a perfect crystal at 0 K allowed into an ideal box at high speed (then the door shut behind it) has a lot of kinetic energy, but no thermal energy, and the system has NO "heat" in that sense (nor does the moving crystal transfer "heat" into it-- just energy). Nor is there a "temperature" to talk about, except the 0 kelvin of the crystal. Likewise, an excited atom has energy available, but it isn't heat, and the atom has no temperature (these are all statistical quantities which take many atoms or states not wone). The energy is there, but not thermalized. Same for an atom that emits ONE photon. The same for two atoms that could react but haven't, etc. The energy to be released, or that has been released, is not thermal energy until it's been spread out among all the particles that can take some of it. Thus you see there are many ways to transfer energy into a system which do NOT involve heating it, but which TURN INTO heat once they get in. Our moving crystal is rather like dropping a book to the ground where it does not bounce: the kinetic energy isn't heat while the book falls, yet it certainly turns into heat (or actually not heat but thermal energy) after the book hits. (This is not a complicated concept, is it?). Once it IS thermalized within the book, good luck in trying to figure out how to get this energy to launch the book upwards again, as high as it was. That's almost the whole of thermodynamics, right there. A bunch of excited atoms (or a cold crystal moving very fast) allowed into a cold ideal box are like the falling book: no temperature and no heat and no thermal energy. But as soon as these atoms impact and use their kinetic energy to emit photons, and the box is filled with a photon gas representing the equilibrium temperature, the atoms are all like the book on the ground, which is filled with the thermal energy of its impact after it hits. You can't reverse this.

Anyway, our rapidly moving crystal, you see, is very much like the book. If the crystal hits the wall and breaks up into gas in the box and the gas takes its portion of the energy until finally the box is filled with a gas at uniform pressure and temperature, now the kinetic energy is ALSO thermal energy. It has been CONVERTED to thermal energy, even though it remains kinetic energy. The extra thing that has been done is to randomize it in direction, and break it up into a distribution of kinetic energies. Assuming a monatomic gas, it remains kinetic energy (STILL) but it's in a form where only a faction of it can be made to do work, unless you have an infinite volume or a reservoir at absolute zero, to help you get it back. Whereas, if the crystal bounced off a spring at the back of the box and came back out at the same speed it went in, and still cold, you could get all of its kinetic energy back (same with the book if it hit a perfect spring and bounced back up to its original height).

But once the energy is thermalized, now you have problems using it. If you let the gas in your "ideal box" expand a cylinder to do work, some heat is always left, and now your gas is in a bigger volume-- you are losing and you see the problem. The entropy of volume expansion has helped you get SOME energy back into useful work, but not all, unless you have infinite volume. Let the atoms out in a beam and you have the same problem of volume. Let them out one-by-one per Maxwell's trapdooor/demon, and they will build up on the other side until you find a volume to get them away again, and you're in the same pickle. And yes, an atom which would like to radiate but is being bombarded by so many photons at the frequency it would like to emit, is like the atoms that are in equilibrium in any object with a temperature-- it is inhibited from emitting, or else it absorbs as easily as it emits, which is the same thing. Now it's just a part of the reservoir holding heat, and its energy represents a bit of heat. If the atom has more energy on average than the reservoir, we are not in equalibrium, and it can emit energy that will eventually be thermalized, and will then be heat. But the atom's extra energy is not "heat" (THERMAL energy) until that process happens.

Ultimately all these problems are related to quantum objects (which have wavefunctions) expanding into a volume. You can't compress them again without doing work, and you can't get work out of them without finding space to expand.

Imagine connecting your box to an equal-sized box "filled" with vacuum, but no photons (it's at 0 K). This is a low temperature reservoir and will accept heat from a connected source at any non-zero temperature, until it fills with photons (which it creates, or are created by the atoms in the walls of the box) and reaches an equal temp. Any reservoir you use to extract heat works the same way, though most have larger heat capacities than vacuum (since they have other ways to store heat energy than just photons in vacuum). They all amount to finding a larger space (volume) which isn't yet filled with the photon "gas" representing the higher temperature of the reservoir you're trying to get work out of, by removing its heat and converting it.SBHarris 18:56, 17 September 2011 (UTC)

"And yes, an atom which would like to radiate but is being bombarded by so many photons at the frequency it would like to emit, is like the atoms that are in equilibrium in any object with a temperature-- it is inhibited from emitting, or else it absorbs as easily as it emits, which is the same thing." Is this what the concept of "vacuum energy" is about?siNkarma86—Expert Sectioneer of Wikipedia
^{86 = 19+9+14 + karma = 19+9+14 + talk} 00:38, 18 September 2011 (UTC)

No, vacuum energy is a theoretical concept which has little to do with this. The vacuum at any non-zero temperature contains photons (as a photon gas that has a meaningful temperature), but the "vacuum energy" is about the energy of vacuum even at absolute zero. What we're talking about with transfer of "heat" to an object is the interaction of normal photons with normal atoms. That even happens with transfer of heat to a vacuum, as the atoms in the container surrounding the vacuum then provide the photons (which are newly made-- photons are NOT conserved). It's just that the photon energy is jumbled up and has a black body spectrum of energies. THEN (and only then) can it be called "heat energy" (thermal energy). The same can be said for the kinetic energy of gas atoms at a given gas temperature, or the energy of motion of atoms in a solid (which includes potential energy) at a given solid temperature. It's not until all this new energy is parcelled out and jumbled up in many different amounts and directions, that it becomes "thermal," and then "temperature" has a meaning. Once more, don't think of all internal energies as thermal energies. Thermal energy, which is called "heat" when it is transfered down a temperature gradient, is a very, very special type of randomized energy, filling all available quantum states, according to certain rules. SBHarris 02:11, 18 September 2011 (UTC)

Heat defined as any energy transfer OTHER than work?! I don't think so!

I haven't checked reference #1, but it must surely be either badly quoted, or itself mistaken. There are many ways of tranferring energy to a body in other ways than due to work, which are NOT heat. A simple example is transferring energy to an object in a microwave oven, or with a laser. Or for that matter, firing a single photon at it. None of these types of radiation are "heat" since they have the wrong spectrum for heat (thermal radiation). This is easy to see with a single photon, since temperature is a statistical quantity, heat requires a gradient, and a single photon can no more have a temperature than can a single atom. Thus, single photons cannot transfer heat, and are not themselves considered heat (though they may be PART of a large number of photons in thermal radiation, which in aggregate, are heat). Thus, this definition is simply bad, and should be removed. Right now, it's stinking up the lede. SBHarris 15:26, 3 October 2011 (UTC)

I am coming to this conversation without knowing its context and I don't know what you mean by reference #1. I am guessing you mean Reif (2000). If I may offer my comments.

The ordinary language word heat is not sharply defined as a technical term in physics, and the ordinary language evolved before the concept of energy was understood. In thermodynamics we are interested in transfer of energy as heat, and then we think of amounts of heat measurable in calories. In theoretical thermodynamics, about 1907 people (Bryan) began to try to think in terms of an axiomatic system, and in 1909 Carathéodory published a rather mathematical axiomatization, which delayed the definition of heat till after statement of the zeroth, first, and second laws. There are people who think this axiomatization is the acme of wisdom, and that anyone who doesn't live by it is a half-wit, and some of them have written textbooks about it. There are others, cited with approval by none less than Clifford Truesdell, who think that: "To deal with the foundations of thermodynamics as though you don't know what temperature, work and heat are is nonsense" (Zemanksy, M.W. (1970) Pure and Applied Chemistry 22: 549-553. From proceedings of the international conference on thermodynamics held at Cardiff UK 1-4 April 1970.), and Zemansky has also written a textbook. Reif himself mixes up statistical mechanics with macroscopic thermodynamics, and there are those who this this is also a sign of great wisdom, and others who think it is muddling. Reif tends to define heat in the manner of Carathéodory. Much, perhaps most, thermodynamic data in practice is measured by calorimetry. Some people try to make out that heat really means entropy, not a form of energy at all, but they are in a minority.

Heat transfer narrowly defined would be only by conduction and by Planck thermal radiation, I suppose. One can define transfer quantities with respect to the source of the energy or with respect to the receiver. Laser transfer could be defined as heat transfer if all it does is to heat the receiver. But if it is received by some device that responds specifically to the wavelength of the laser light, and transfers the energy directly into some form of potential energy, than it could be defined as transfer as work. There is a very good case to be made that it is more convenient to distinguish work transfer from heat transfer by how the transferred energy is received, not how it is sent; but the other way can also be defended. According to Adkins 1968/1975 page 34: "The distinction between heat and work is not always clear-cut ..." ((((Let us not even mention "dark energy" here!))))

My personal preference would be to just delete from the lead the sentence: "In this description, it is an energy transfer to a body in any other way than due to work.[1]" but there are those who would have me shot at dawn every morning for a year for such temerity.Chjoaygame (talk) 17:58, 3 October 2011 (UTC)

I disagree, the definition given by Reif is the only correct fundamental definition that is applicable to any arbitrary system. It's only when you consider systems close to thermal equilibrium that you find the relations between temperature gradient and heat flow, but in general, a system can be very far from thermal equilibrium and so you want a definition of heat that works generally. This is particularly important since a heat flow already implies a departure from thermal equilibrium.

E.g., suppose you start with two systems, each at thermal equilibrium but at different temperatures and you bring them into thermal contact via an insulator so that the heat only flows very slowly. Then that heat flow will also perturb the state of each individual system away from thermal equilibrium. You can make that departure from thermal equilirium arbitrary small by increasing the insulaton and let hte heat flow rate become arbitrary small, but you can't make it zero while still have a non zero heat flow. This is why books such as the one by Reif that treat the subject rigorously wil always be careful when applying thermodynamics, always making it explicit where changes happen "quasistatically" etc. etc.

What this means in practice is that to derive the heat conduction coefficient from first principles using the differential cross section of molecules, you must treat a gas in a non-equilibrium situation, the velocity distribution is then perturbed from the Maxwell distribution. So, you don't just have a Maxwell distribution a position dependent temperature, you have to include addition terms. So, tying heat down to temperature differences is from a rigorous theoretical point fo view problematic, because heat flow happens precisely when the concept of temperature (defined locally) starts to break down.

Then considering a completely general system, all you can do is specify what degrees of freedom you want to describe statistically and what you are going to keep track of explicitely. This defines the distincton between heat and work, which is thus in principle competely arbitrary. In practical settings, you encounter this freedom when you decide where to put you system boundary. E.g. in the free expansion experiment, you can decide to treat the two chambers separated by the piston as one system and then the system doesn't perform work and there is no heat flow into the system. You can also divide the system into two parts and keep track of the position of the pistion and some details of the gas flow as it approaches its new equibrium. Then, you can say that the gas performs work by accelerating the piston and look at heat generated from friction etc. etc. Count Iblis (talk) 18:03, 3 October 2011 (UTC)

I will agree that if you look too "closely" (too microscopically) at heat (flow) it begins to depart from classical heat, just as when you look too closely at a small bit of material at a given (stable) temperature, it begins to depart from even having a classical temperature! So what? Such bad things happen when we look at small sample-sizes for concepts which depend on large sample sizes, since they are statistical. That's a bad thing to do, and not allowed. It's rather like saying that "wetness" is not what we think it is, because when you go down to the level of individual water molecules, you can't "find" it. And thus suggest that wetness must be re-defined, or perhaps is unreal/arbitrary. To which I can only say: "no"! You're merely looking for a system property at the wrong level of complexity, and that's fruitless to do as a hobby. It will give you nonsense for heat, gas pressure, gas temperature, and all sorts of other stuff. None of them really "arbitrary"-- just more uncertain at small scales. SBHarris 18:41, 3 October 2011 (UTC)

• Maybe not. Let me also formally record my support for deletion of this sentence. Another example: if I attach a pair of wires to a block of metal and run an electrical current through it, it's going to warm. I've certainly "heated it up," meaning loosely and vulgarly that I've raised its thermal energy and temperature. But just because the energy has appeared as heat in the block, doesn't mean it was necessarily transferred as heat. But it wasn't transferred as work, either: I ran a current through the thing, but I did no work (mechanical work, which is the only kind of work thermodynamics recognizes) upon it. So, yes, the lede definition is bad. SBHarris 18:09, 3 October 2011 (UTC)

There is no problem here, work can be dissiptated into heat. You can keep track of an external parameter of one system, the energy change associated with the change in that is work performed by that system, while when that energy flows into another system, it's not attributed to any external parameters of that system, so there is counted as heat.

This is a consistent framework of describing things, actually more consistent than the way engineering and chemistry books do it (what they do assumes that you are close to thermal equilibrium and that it is always obvious what the external parameters are). Count Iblis (talk) 18:29, 3 October 2011 (UTC)

There is a problem here! Just because some energy winds up as heat, does not permit us to define it as heat. I can put a metal ring in a changing magnetic field, and it will heat inductively. Heat is produced, but is there any heat transfer that a thermodynamicist would recognize? No. Nor work, either. If I didn't allow the induced current to dissipate as heat, I could have used it as electricity, as in the secondary of a transformer. In that case, clearly most of it doesn't become heat, and never was heat. SBHarris 18:41, 3 October 2011 (UTC)

As I mentioned, some authors believe that the Carathéodory way is the only right way to define heat, and others think it is a travesty of reason. Count Iblis seems to believe it is the only correct way, and quotes a pedagogical textbook that supports that view. But other authoritative books disagree. Besides seeming in this respect to be a staunch Carathéodorist, Count Iblis seems to have little regard for a distinction between the macroscopic viewpoint that is often considered to characterize thermodynamics and the microscopic viewpoint that characterizes statistical thermodynamics, though Carathéodory's work itself is strictly macroscopic with no microscopic element. For one of many possible examples, Pippard, A.B. (1957/1964), Elements of Classical Thermodynamics for Advanced Students of Physics, Cambridge University Press, writes: "...classical thermdynamics ... takes no account of the atomic constitution of matter..." There are several or even many defensible points of view. Is it Wikipedia policy to dictate the one and only correct viewpoint?

As I mentioned, the energy tranferred can be classified according to how it is sent to the system or to how it is received by the system. There is a case that it should be classified according to how it is received by the system. It is not out of the question to say that the relevant parcel of energy leaves the environment as work and enters the system as heat. This question is analyzed at Gislason, E.A., Craig, N.C., Cementing the foundations of thermodynamics: Comparison of system-based and surroundings-based definitions of work and heat, J. Chem. Thermodynamics 37 (2005) 954–966. I think if one wanted clarity and consistency, it would be necessary to say that there are two ways of doing it and to state them both, and give reasons for choosing one or the other in particular circumstances.Chjoaygame (talk) 20:51, 3 October 2011 (UTC)

But the statement in the lede is wrong if the energy delivered to a system never is received as heat, for then the energy never-ever is heat, period. As when it is in the form of electric or chem potentials-- as when you charge up a capacitor, or a battery, it isn't heat or work. Or (same thing as a charge) remove burned fuel and add unburned fuel. You add chemical potential to a system that way, like changing a spent battery cell in a flashlight rather than re-charging one in-place. Thus, the statement is wrong even for classical (non microscopic) thermodynamics. There just isn't any heat if there's no heat, dangit! You can look all you like, but no classical heat means no heat of any type, period. SBHarris 22:26, 3 October 2011 (UTC)

Work doesn't have to be mechanical work. The work done by a system is defined as the decrease in the internal energy of a system due to the change of its external parameters. If we have one external parameter x, then for fixed x we have an isolated system which has energy levels that depend on x. There are then energy eigenstates |n,x> with energy eigenvalues E_n(x). Then it follows from the quantum adiabatic theorem, that in the limit of an infinitely slow change of x, the system in state |n,x> won't jump to another state |n',x> with n' different from n. This means that the work done is by the system is -dE_n(x)/dx dx. One then defines X = -dE_n(x)/dx to be the generalized force conjugate to x, so the work done can be written as X dx in the quasistatic case. The change in internal energy due to work done is thus - X dx.

If we have many external parameters x_i, the quasistatic work is sum over i of X_i dx_i. Then the definition of heat as "what is left after taking into account work", leads to dE = dQ - dW. And one can then show, using this definition that the change in entropy in the quasistatic case is given by dS = dQ/T. So dS does not depend on the changes in the external parameters, see here for a derivation. What you then get is that:

$dE=TdS-\sum _{i}X_{i}dx_{i}$

But now comes a very important step. The internal energy of the system in thermal equilibrium only depends on S and the external parameters x_i. This means that while the above equation was obtained by considering only quasistatic changes, it is valid in general. So, if you rapidly change the external parameters, then the work is no longer given by the X dx term, but also the heat is not equal to T dS, but the above equation still holds (if you wait until the system has settled down into thermal equilibrium after the rapid change). X dx is then more than the performed work, you can then interpret the difference as "work being dissipated into heat" leading to the extra entropy increase. Count Iblis (talk) 23:49, 3 October 2011 (UTC)

I'm not talking about systems where there is an entropy increase. I'm talking about systems where you add energy with NO entropy increase. I think you want to define all such energy increases (like charging a capacitor or simply adding cold mass) as doing "work" on the system. However, this is defining "work" in a way that isn't very helpful, simply as any (non heat) energy added to a system. Then "heat" becomes "any non-work energy" added to a system, as a matter of syllogism. However, in the process, we've defined work in way that isn't common in physics. Physics, commonly, by the word "work" means "mechanical work." That's how we define work (physics) on Wikipedia, too. It's mechanical work. SBHarris 00:53, 4 October 2011 (UTC)

It is indeed a syllogism, see e.g. also Work (thermodynamics). There is just no other way to define heat in general. It's ultimately the energy that is transferred to all those degrees of freedom that we don't keep track of and describe statistically. In principle you can decide to keep track of every molecule in the system, then you have as many external parameters as there are degrees of freedom in the system. Then all energy transfer counts as work, there is then no heat and the entropy is always equal to zero.

You only obtain nonzero heat and entropy when you give a coarse grained description of the system, keeping only a few external parameters. This coarse graining is an essential step, you need to define the Omega function that counts the number of states by choosing a finite energy resolution delta E. Then Omega(E) is the number of states with energy between E and E + delta E. The entropy is defined as k Log[Omega(E)]. This thus depends on delta E, but this dependence on delta E is then completely negligible. You can interpret the entropy is the amount of information you need to give to specify exactly which state the system is in, given that you know that the energy is between E and E + delta E. That amount of information is proportional to the number of degrees of freedom of the system, say 10^23 bits, and changing Delta E by a factor of 2 would add 1 bit to this.

But if you were to choose Delta E so small that there is only one energy eigenstate within the energy interval then S becomes zero. All the information about the system's state is now in the thermodynamic specification of the system's energy. Count Iblis (talk) 02:47, 4 October 2011 (UTC)

Okay, then in the lede you need to explicitly mention that you're talking about work (thermodynamics) and not work (physics), which is mechanical work. Work (thermodynmics) is the mere CAPACITY to do mechanical work, whether it is done or not. So it could involve electric current or fuel transfer. As such, it is not much different from thermodynamic free energy (in fact, if you're going to use "thermodynamic work" as a concept, what is the point of inventing "free energy" as a concept?). In any case, I don't see that this definition of work is very helpful here in the definition of heat, since one needs to understand heat first, in order to obtain a figure for this definition of work. SBHarris 16:46, 4 October 2011 (UTC)

Since, in the second law of thermodynamics and all matters to do with heat engines, heat is measured by temperature which is related to energy by the Boltzmann constant. By the general definition of energy it can be converted to (or from) heat e.g. when water turns to steam (or the other way round). This can be done by burning fuel (adding heat directly and irreversably) or {doing work}/{letting work be done} e.g. with pistons and cylinders compressing (or expanding) the steam. By 'burning fuel' you can change water into steam. But you cannot turn water into steam by doing work on it directly. However you can turn steam into water by compressing it with a piston.

To understand heat you must first understand the relation between energy and temperature, which is the Boltzmann constant. Then you can move on to entropy and some of the other more difficult concepts associated with molecular energy. At present nobody seems to want to discuss the Boltzmann constant! This is a shame because the article is about heat and heat is about the Boltzmann constant. --Damorbel (talk) 18:28, 4 October 2011 (UTC)

Heat and the Boltzmann Question

When do we agree that

1/ the Boltzmann constant is the key relation between energy and temperature?

2/ When do we agree that Heat is measured by temperature?

3/ When do we agree that thermal eqilibrium reqires either uniform temperature distribution or a Maxwell-Boltzmann distribution of particle energy? --Damorbel (talk) 19:30, 11 October 2011 (UTC)

In answer to the above:

2). It isn't. Heat is a quantity of energy, and is measured in Joules.

3). It doesn't. In general the equilibrium distribution depends on the nature of the partition function. Besides, it's irrelevant to an article on Heat.

1). Arguably, the key relation between temperature and energy is the equation:

${\frac {1}{T}}={\frac {\mathrm {d} S}{\mathrm {d} E}}$

which can if one wishes be treated in entirely macroscopic terms, without any reference to the Boltzmann constant.

People have gone through all this with you above. Is it that you can't read, or that you simply don't have the capacity to engage with what people have been trying to explain to you? Jheald (talk) 22:10, 11 October 2011 (UTC)

Let me add that there isn't anything magical about the Boltzmann constant, which is not some mysterious dimensionless constant like the fine-structure constant, but instead is once of those constants like c, which has a value that merely reflects that fact that it connects two different measuring scales (just as the value of c connects scales of length and time). In the case of k and R, they connect the SCALES of temperature and energy that we choose to use, and we could choose scales where k and R were 1 for any system. One more problem is that T reflects only the part of thermal energy that is kinetic, not the total that is thermal energy (these are not the same except in systems like an ideal gas where all thermal energy is kinetic energy). So there is one more number that separates E and T which to do with degrees of freedom in a system storing heat. This number is pretty small, and varies from about 1.5 to 3 per particle.

Anyway, all this prevents thermal energy from being measured as temperature, even if we decided to scale temperature so we didn't need a constant of proportionality. Temperature is proportional to energy/particle or energy/mole, so we COULD have chosen to measure temperature in "joules/per mole." or "joules/particle" But those are not quite the same units as heat, but rather heat capacity. And we'd still be stuck with the degrees of freedom thing, even so. So we've introduced k and R = k/N to take care of "per unit" problems (converting heat capacity to thermal energy, so that every temperature has an energy), and we still need a pure number to take care of the degree of freedom problem. And when that is done, here we are. But energy and temperature can't be equated without taking care of those two problems. SBHarris 23:05, 11 October 2011 (UTC) )

The fundamental issue here though is that this is an article on heat not temperature. Heat is not temperature. It is not even the same concept as thermal energy. Until User:Damorbel grasps that fundamental point any more sophisticated discussion here is besides the point. Jheald (talk) 23:12, 11 October 2011 (UTC)

"Heat is not temperature" Then what does temperature measure? Further, why is the measure of temperature (and by extension entropy) the central parameter of the 2nd Law of thermodynamics? --Damorbel (talk) 18:03, 8 December 2011 (UTC)

Two bodies at the same temperature may contain different amounts of heat. In one case, the two bodies may be made of the same stuff but have different sizes. In another case, the bodies may simply be made of different stuff (for example, copper and water). Heat is the amount of energy. Temperature a measure of that energy, but also depends on the specific heat of the body. In addition, a change in heat can cause a substance to change state (e.g., solid to liquid) with almost no change in temperature. Q Science (talk) 07:31, 9 December 2011 (UTC)

Q Science, I see it that way also. Jheald is using E instead of Q in his formula above which is incorrect, Q is the thermal energy which is related to temperature; as you point out state change involves change in energy (E) without change in temperature, thus temperature and energy are not related in such a simple way. --Damorbel (talk) 08:06, 9 December 2011 (UTC)

Further on this matter of Q and E. Jheald is not alone in using E, Frederick Reif in his oft quoted book 'Statistical and Thermal Physics' uses E in his definition of entropy on p99 (his formula 3.3.11 is the same as Jheald's). When discussin energy E a consistent result can only be obtained if you define what energy you are considering, when it is temperature it has to be molecular kinetic energy (heat). As an example, suppose your gas was hydrogen H2; not much different from nitrogen you might think, but suppose the conditions were such that the hydrogen fused to helium, or it might be uranium fissioning. In these cases the energy is quite different from the simple Q of the thermal case. Q is still very important but the fusion and fission energy becomes very important also and absolutely cannot be ignored. Even H2 and N2 are not that simple, they are diatomic and, at a given temperature, contain more energy per particle than monatomic gases such as argon and helium.--Damorbel (talk) 08:02, 10 December 2011 (UTC)

Response to Darmobel. It can save time to define one's intentions in ordinary mathematical and thermodynamic terms. The equation posted by JHeald can also be written

$\left.{\frac {1}{T}}\right|_{(E,V)}=\left({\frac {\partial S}{\partial E}}\right)_{V}=\left.{\frac {\partial S}{\partial E}}\right|_{(E,V)}$

It refers to the so-called 'entropy representation' of a closed system for which the fundamental equation is $S=S(E,V)$ where $E$ denotes the internal energy (which is denoted by the symbol $U$ in some other discussions), and where $S(E,V)$ is a function of $(E,V)$ that gives the dependent variable entropy $S$ for the independent variables $E$ and $V$. In this case, all the variables are variables of state and the equation defines the temperature $T$, which is a state variable, a function $T(E,V)$ also of the independent variables $E$ and $V$.

Formulas about $Q$ and $S$ have a different meaning. For example, they might refer to a process of volume change from $V_{1}$ to $V_{2}$, that goes so slowly that it may considered to be reversible. The symbol $Q$ then represents the accumulated amount of heat transferred to the system during the process, which we specify in terms of $Q$ and $T$ as functions of $V$. Then we can find the entropy change for the reversible process by calculating

$S_{2}-S_{1}=\int _{V_{1}}^{V_{2}}{\frac {1}{T(V)}}{\frac {\mathrm {d} Q(V)}{\mathrm {d} V}}\mathrm {d} V$

Chjoaygame (talk) 02:10, 11 December 2011 (UTC)

Chjoaygame,I suggest your explanation illustrates my argument. You refer to E as an independent variable which is fine but if you put it in an equation where it is a function of T there is only one relation that can exist between E and T and it is the exchange of momentum (energy?) by random elastic collision of particles as described by the Maxwell-Boltzmann distribution. I suggest it is fair and correct to descibe this energy as Q, the thermal energy of the system. In any 'real' system there other other energies present, chemical, gravitational etc., so the total system energy will be above the thermal energy. However, as soon as you introduce temperature the system must be in either in equilibrium or at least a steady state i.e. the is no energy exchange taking place other than the random collision of particles. The physics has to take precedence over the mathematics! --Damorbel (talk) 10:11, 12 December 2011 (UTC)

focusing again on the article

For the sake of the article, we may try to collect our thoughts.

Heat is a word of the ordinary language, but also a term of art for engineers, chemists and physicists. I am inclined to say that chemists and physicists should accept the thermodynamic view of heat, so that would make heat a term of art for chemists and physicists. I am very reluctant to accept the idea that heat is fundamentally different for engineers and physicists, though it is likely that there are various usages that are more or less abuses of language occur in physics and engineering. I am in favour of avoiding abuses of language in this article, as far as practical, and of indicating in the article that abuses of language do occur, and perhaps describing them.

I am unhappy with the phrase "In engineering and [layman] [laymen] language". The usual English would be "layman's language" or "laymen's language" or "lay terms". I would prefer to write of ordinary language and technical terms or terms of art in various disciplines. I am inclined to put physicists' and engineers' usages together as distinct from ordinary language usages, rather then putting laymen's and engineers' usages together as distinct from physicists'.

In physics it is traditional to speak of heating or doing work on bodies, not "objects".

Generally speaking, thermal energy is generated by processes or pathways, and only in particular cases by "methods".

In the parts of the article that concern thermodynamics, strict thermodynamic presentation should be adhered to, and looser forms of presentation should be thoroughly avoided.

Heat transfer by friction and by viscosity is not fully determined by the overall temperatures of the systems.

No reliable source has been provided for the removal of the idea of heat transfer by radiation from the lead. This will remain the case, because no reliable source can exist for its absence. The policy of reliable sourcing does not have exemptions, see WP:IRS. In our area, the custom seems to be to ignore the policy of reliable sourcing for the leads of articles, but that is not properly justifiable. Disciplines with advanced theories and many points of view need particular care and sober and judicious selectivity for identification of reliable sources, and the mere fact that a source is reliable for one question does not make it reliable for others.

Thermal radiation is a form of heat transfer in thermodynamics. Reliable sources for radiant heat in thermodynamics are

• Planck, M. (1914), The Theory of Heat Radiation, second edition translated by Masius, M., P. Blakiston's Son & Co., Philadelphia.
• Partington, J.R. (1949), An Advanced Treatise on Physical Chemistry, volume 1, Fundamental Principles. The Properties of Gases, Longman's Green and Co., London, Chapter VI, Part B, Section 4, pages 466–468.
• Guggenheim, E.A. (1949/1967), Thermodynamics. An Advanced Treatment of Chemists and Physicists, North-Holland Publishing, Amsterdam, LC 67–20003, Chapter 12, pages 357–361.
• Landsberg, P.T. (1961), Thermodynamics with Quantum Statistical Illustrations, Interscience Publishers, New York, pages 258–305.
• Mandel, L., Wolf, E. (1995), Optical Coherence and Quantum Optics, Cambridge University Press, Cambridge UK, ISBN 0–521–41711–2, Chapter 13, pages 659–682.
• Kondepudi, D. (2008), Introduction to Modern Thermodynamics, Wiley, Chichester UK, ISBN 978–0–470–01598–8, page 53.

Sources that might be reliable for other matters, but omit radiant heat for thermodynamics, are not reliable as sources for a proposed view that thermodynamics does not include radiant heat transfer.

It would be right to want to make the finer point that wavelength-specified radiation has a temperature, and if its temperature exceeds that of the receiving body it can be accepted as work rather than as heat, but the simple removal of the idea from the lead does not make this clear. It would take some effort to make it clear.

Without that effort having been made, the simple idea of heat transfer by thermal radiation should stand in the lead.Chjoaygame (talk) 23:33, 13 December 2011 (UTC)

Sorry, but I do not understand what you mean when you write "wavelength-specified radiation has a temperature". Analysing or exlanation of temperature or heat on the basis of wave functions can only result in failure - should it not be photons? --Damorbel (talk) 21:13, 14 December 2011 (UTC)

The temperature of a monochromatic pencil of radiation is defined by Planck 1914, and by Landsberg 1961 page 295. Planck is perhaps easier to understand. It is not necessary to talk of photons for this purpose.Chjoaygame (talk) 21:28, 14 December 2011 (UTC)

This link lets you search a translation of Planck's "The theory of heat radiation" Planck makes it clear many times that the only radiation with a characteristic temperature is 'black' radiation, e.g.§70 (p68). If you still feel that an (unaltered) monochromatic beam can have a temperaturethen I would like to have a direct quotation.

Strictly speaking a monochromatic beam has zero bandwidth and thus can contain no energy, neither can it be a 'beam', it is a cone (§16, p14 & §18, p16).

In §52 (p44) Planck explains how arbitrary radiation (thus a monochromatic 'beam') into blackbody radiation by inserting a microscopic carbon particle into an otherwise perfectly reflecting cavity.

On p89 of §93 Planck writes "monochromatic radiation, which is uniform in all directions and has a definite energy density u, has also a definite temperature given by (117), and, among all conceivable distributions of energy, the normal one is characterized by the fact that the radiations of all frequencies have the same temperature." From this you may conclude that "monochromatic light" has a temperature; Planck's argument is perfectly logical in that each individual frequency component dv of black radiation has the same entropy, however these frequency components each carry only a fraction dv of the total black body radiation energy so the time average of isolated dv fractions will yield a lower temperature than the blackbody temperature; this is clearly to be seen in the way the Sun at 5780K can only warm the Earth to 279K despite the (mean) photon energy of sunlight coming from a 5780K source.--Damorbel (talk) 18:58, 15 December 2011 (UTC)

I do not have time to argue with you about this. Planck, if you read him carefully, talks mostly of pencils, as I noted above, not of "cones" as you say above. Your comment about the sun heating the earth is not right; you need to think it over a bit more. I suggest you read Planck a bit more, and think it over a bit more.Chjoaygame (talk) 19:43, 15 December 2011 (UTC)

Chjoaygame, you wrote "I do not have time to argue with you about this". Then why are you writing in the 'talk' of this article? I happen to think the article has shortcomings and I would like to improve it. I would prefer not to think I was wasting my time.

Indeed Planck does refer to 'pencils', well over 100 times but the first mention is "Every point of dr will then be the vertex of a pencil of rays diverging in all directions" I suggest you be a little tolerant of the translation, pencils don't really diverge! (Oh well, at the point then.) --Damorbel (talk) 20:03, 15 December 2011 (UTC)