Talk:Enthalpy

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Query

I would really like to know what enthalpy means: help anyone? —Preceding unsigned comment added by 62.252.0.5 (talk) 00:42, January 9, 2005 (UTC)

I myself am just an undergraduate student at the University of Illinois. It appears the word comes from the greek language and means "to warm", or "to heat". A possible meaning of the word could then be expressed as the amount of warmth (which is amount of energy) in a room at a certain temperature and pressure.

For example: In a house each room could be assumed to be at the same pressure (atmospheric pressure). However, the garage would be at a lower temperature than the living room. The amount of enthalpy would be less in the garage (lower temp = less warmth,energy) than in the living room (higher temp = more warmth,energy). —Preceding unsigned comment added by 12.203.115.91 (talk) 16:52, January 31, 2005 (UTC)

I've believed for many years that enthalpy is an antonym of entropy - to have a tendency towards a more ordered situation (as opposed to entropy), opposite that of moving towards disorder (entropy). Checking with a dictionary, this appears to be a consistent view with the meanings of the words in one particular sense - that of enthalpy being defined as a measure of the ability to perform mechanical work, which will only increase as a function of order. Likewise, the definition of entropy in one context is as a measure denoting the disorder of the system being considered. The mathematical discussion of entropy / enthalpy reflect that in their formulas, but it is left to infer those definitions as a consequence to the understanding of the reader. Perhaps a short discussion of the meaning of the words in a more abstract sense would be beneficial to assist other readers? —Preceding unsigned comment added by 24.255.211.121 (talk) 08:02, July 5, 2005 (UTC)

The following was moved here from a section called "Agreed....." by Riick (talk) 05:01, 2 August 2010 (UTC) [reply]

== Agreed..... ==

I strongly agree. I'm not at all trying to divert this entry from hard science, but I think that if enthalpy is not a tendency towards order, there must be another term for the tendency towards order. I've been wondering about this for quite some time, maybe someone has an answer?

I have this feeling that as hard as entropy is to quantify, so must be enthalpy...... How on earth does a tree grow in the forest when the entropic nature of the universe, as we've been taught to believe, creates an inevitable pull towards a state of unified disorder, the decay and rot of natural matter.

There is a certain duality in forces here, and it just seems to me that enthalpy is the natural opposite..... I just don't know. And I can't seem to find anything defining or discussing enthalpy as a tendency towards order, save for a few sparse references in a few early-mid 20th century physics volumes. And in such cases, there was barely enough to go on.

Perhaps there is a nice Professor of Physic roaming the wikii that might enlighten us. —Preceding unsigned comment added by 69.221.174.207 (talk) 07:15, July 12, 2005 (UTC)

deolijn: umm just a question. entthalpy is a move towards order but when you mention a tree please understand it is a solid and thus different rules apply —Preceding unsigned comment added by 58.162.26.205 (talk) 12:35, July 21, 2006 (UTC)

Newbie: I'm no scholar, but i've always used "orthotropy" to be the antonym of entropy. If entropy means "turning in" or "in-turning" (like the way food turns?), then "orthotropy" would mean "right-" or "correct-turning". —Preceding unsigned comment added by 66.236.159.171 (talk) 21:47, July 30, 2005 (UTC)

You might like to take a look at Second law of thermodynamics for more of a discussion: enthalpy is only one of the factors involved. Physchim62 (talk) 15:09, 9 December 2005 (UTC)[reply]

Unregistered Roger: If you're trying to relate entropy to enthalpy then I'm afraid you're going nowhere. Entropy is a statistical measure and indicated what is the most likely state(ie.the one with the highest entropy) and the direction of the change in state. Enthalpy is a measure of energy and is not a fundamental property, but a convenient compound of 'internal energy' and 'pressure x volume'. It turns out that in a system where pressure is constant or where we are looking at the steady flow of a fluid through a system, then the change in enthalpy gives us a measure of the amount of energy that must flow into that system. —Preceding unsigned comment added by 137.205.146.206 (talk) 15:44, June 22, 2006 (UTC)

Freshman: I'd like to know whose P in the enthalpy equation refer to. The system's or the surrounding's pressure. Some articles listed it as surrounding's. This wikipedia article listed it as the system's. Please clarify the confusion. —Preceding unsigned comment added by 128.253.197.151 (talk) 05:08, October 23, 2006 (UTC)

Not trying to question hard science but seriously, a young and up and comcming scientist faced with this as the opening sentance. "In thermodynamics and molecular chemistry, the enthalpy (denoted as H, h, or rarely as χ) is a quotient or description of thermodynamic potential of a system, which can be used to calculate the heat transfer during a quasistatic process taking place in a closed thermodynamic system under constant pressure. " As one of my old teachers said to me. "If I hear another bloody quadrillion-syllabic word out of your mouth, detention!!" how about a nice simple intro?? Then you can lose the young whippersnapper later in the page. —Preceding unsigned comment added by 92.238.94.146 (talk) 02:39, 19 April 2009 (UTC)[reply]

(In reply to the one ~~below~~ above). Guys, enthalpy is NOT the antonym of entropy. If you have taken a class on thermodynamics, it is obvious that enthalpy is the sum of U and PV (i.e. H = U + PV). There is a great possibility that you are mistaking the concept of enthalpy with Gibbs Free Energy (G), which is G = H - TS. Remember that the equality sign here represents the system at equilibrium, otherwise you are going to have a <= sign there. Thank you and please refer back to your thermodynamics notes before writing on any science entry.

Guys, the tendency towards order is obviously the negative of entropy (-S). Thanks!

Guys, we can use Maxwell's Equation for it. Ouch man! —Preceding unsigned comment added by 168.122.13.29 (talk) 22:48, 27 March 2012‎

Redirect

1. ) "Standard enthalpy change of denaturation" should link to a page titled "Standard enthalpy change of denaturation," not, as it now does, to the page on "Denaturation." This does not define ∆H(denat). Or, if "denaturation" is the preferred link, only that word should be highlighted. (12sep09) —Preceding unsigned comment added by Schmin (talk • contribs) 20:56, September 12, 2009 (UTC)

2.) Would it be reasonable to redirect heat of reaction here? I don't know enough about the topic, but it's a request article. A redirect here would seem sufficient to me, but can someone more knowledge inform? Would be greatly appreciated! Agentsoo 18:18, 14 August 2005 (UTC)[reply]

RE: redirect

redirect from enthalpy to heat of reaction? ... enthalpy has uses and meaning outside chemistry like the thermodynamic one treated here so while a link might be appropriate, it shouldn't redirect automatically —Preceding unsigned comment added by 137.99.21.199 (talk) 18:57, October 25, 2005 (UTC)

Heat of reaction now redirects to Standard enthalpy chnage of reaction. Physchim62 (talk) 15:09, 9 December 2005 (UTC)[reply]

No. This would not be reasonable. The redirect to Standard enthalpy change of reaction is appropriate for now, but ideally there would be a separate Heat of reaction article with the alternate name Enthalpy of reaction to reflect usage under non-standard conditions. I am removing the redirect tag. Flying Jazz 03:08, 13 February 2006 (UTC)[reply]

contradiction? - is dQ = dH in isobaric chemical reactions?

On the one hand it is written here:

It is seen that, not only must the Vdp term be set to zero by requiring the pressures of the initial and final states to be the same, but the

\mu _{i}dN_{i}

terms must be zero as well, by requiring that the particle numbers remain unchanged.

But on the other hand, the situation in which the use of dH=dQ is the most important, is in chemical reactions - where the sign of dH decides whether the reaction is exothermic or endothermic. But on chemical reacions $\mu _{i}dN_{i}\neq 0$ , and usually (under constant T and P) $\ \sum _{i}\mu _{i}dN_{i}=dG$ , which in spontanous reactions is smaller then 0.

So what do I miss? eman 15:56, 28 June 2006 (UTC)[reply]

The last few paragraphs of the "some useful relationships" section may be confusing you. I think the writer has the goal of showing the situations where the first and second law combine to result in the requirement that dH must be less than or equal to zero. Next, the writer includes the posibility of chemical reactions, but only to once again show the situations where the first and second law combine to result in the requirement that dH must be less than or equal to zero. I think that what the writer means by "interpretation of the enthalpy" when writing "any further generalization will add even more terms whose extensive differential term must be set to zero in order for the interpretation of the enthalpy to hold" is an interpretation of enthalpy as a potential whose derivative must be less than or equal to zero.

I think that the writer is trying to say is "The second law means that dH must be less than or equal to zero if entropy, pressure, and the number and types of particles remain unchanged." In other words, you're not missing anything because when there's a chemical reaction, the writer's "interpretation of the enthalpy" no longer applies. By the way, I think you meant "exo-" and "endothermic." I'll try to edit this in a couple days. Flying Jazz 03:29, 29 June 2006 (UTC)[reply]

Yes, of course I meant exothermic. I corrected myself above.

My main problem is that I know that chemists use the citeria of the sign of ΔH to determine whether a reaction is exothermic or endothermic, and now I can't see how this can be justified? So could it be that all the chemists, and all the chemistry books are wrong?!

eman 15:21, 29 June 2006 (UTC)[reply]

Enthalpy is just heat content, so the chemists are correct by definition. I guess I don't understand why you think they are incorrect. If you take some of the equations you wrote above and combine them with some of the equations in the text, you'll get the equation ΔG = ΔH - TΔS which should be familiar and might explain some of what's going on. Believe the textbooks. The section of the article that you're confused about should be ignored and will be removed soon. Flying Jazz 03:45, 1 July 2006 (UTC)[reply]

What bothers me is that for cheical reactions ΔG is usually not zero (for spontanous reactions it should be negative). So

\Delta G\neq 0

means that

\Delta H-T\Delta S\neq 0

, thus

\Delta H\neq T\Delta S

, . eman 10:41, 1 July 2006 (UTC)[reply]

Yes. The equations that use the equality only apply for reversible processes (where delta G = 0). This is the case at equilibrium. A spontaneous chemical reaction isn't at equilibrium. After the concentrations change or other reaction conditions change so that equilibrium is achieved, the equalities hold. Flying Jazz 12:25, 1 July 2006 (UTC)[reply]

Still it doesn't solve my original problem about the

\ \mu _{i}dN_{i}

in dH. The fact that the Enthalpy is also named "heat content" doesn't help. Nomeclature can be wrong, and has only symbolic importance. eman 11:17, 2 July 2006 (UTC)[reply]

OK, I think I understand what you're asking now, and it's similar to the question you asked on Talk:Internal energy with a similar answer, and it's not a simple answer and I might not explain it fully because I don't understand it 100% myself, but I'll try. Just as there are many types of internal energies depending on whether no $\ \mu _{i}$ , some $\ \mu _{i}$ or all $\ \mu _{i}$ are natural variables, there are also many enthalpies for the same reason. For a "black box with some unknown amount of unknown reactions" inside:

dH=TdS+VdP+\sum _{i}\mu _{i}dN_{i}\,

and all the natural variables (S, P, and all Ni) are independant of each other. But when you have a "heat of reaction" for a defined chemical reaction, all of the Ni are no longer independant of each other. They are related by stoichiometry, so one of the $\ \mu _{i}$ must be treated as a natural variable instead of a conjugate variable and subtracted out as discussed here and here. The nomenclature has been ambiguous for a long time, and maybe Alberty's IUPAC paper will change that in the future. Flying Jazz 16:51, 2 July 2006 (UTC)[reply]

I want to know the enthalpy —Preceding unsigned comment added by 80.191.144.10 (talk) 07:57, November 11, 2006 (UTC)

Article has become messy

The article needs to be re-invigorated by explaining alongside definitions of thermodynamics. I was actually slowed down by downloading an image, and the article looked horrendous when the MathTeX markup wasn't on the screen. There needs to be some definate walkthroughs of the actual processes, rather than just explaining different aspects of enthalpy. This is the wikipedia, not a textbook! J O R D A N ^[talk ] 16:33, 20 February 2007 (UTC)[reply]

ΔH

Im not sure if this has been mentioned, but in the intro but it describes enthalpy as either H or ΔH, Δ actually denotes the change in a value, so ΔH is actually the change of enthalpy, not the enthalpy itself. —Preceding unsigned comment added by 81.178.106.177 (talk) 22:34, April 22, 2007 (UTC)

yeah, so change it. —Preceding unsigned comment added by 86.27.184.240 (talk) 12:57, 15 February 2008 (UTC)[reply]

The product of an exothermic reaction will be warm to the touch.

Is this necessarily true? Can an exothermic reaction not occur at -200 degrees, and be very very cold to the touch?

Ariel Hoffman —Preceding unsigned comment added by 87.69.68.65 (talk) 17:36, 31 December 2007 (UTC)[reply]

You are right, this is not always true. The statement about "being warm to the touch" is only meaningful if we assume that the reaction started at room temperature and the reaction is fast enough so that the products increase the temperature noticeably before losing the heat to their environment. Of course, if the reaction starts at a high temperature, it can feel warm to the touch whether the reaction is endothermic, exothermic, or not even happening at all! --Itub (talk) 11:06, 7 January 2008 (UTC)[reply]

Molar enthalpy

The requested natural sciences page included a request for "molar enthalpy." I redirected here, as per "standard enthalpy" - should we add a short definition of molar enthalpy to the "definitions" section? Coanda-1910 (talk) 02:27, 6 April 2008 (UTC)[reply]

Useless for the layman

There's an old quote that goes "you're so high-minded that you're of no earthly good." That's what I'm seeing out of a lot of the scientific/technical articles on Wikipedia, including this one. The articles are written on a level that is not appropriate for the layman to quickly grasp the concept. Instead, they drift into lofty and even high-theoretical issues without ever giving a firm grounding.

In this case, I used to know what enthalpy was. It was a term I used a lot back in my nuclear power days. However, I've been out of that industry for a while, and found my knowledge was getting a bit rusty. I took a look at this article, and now I still don't know what it is. That's pretty bad when someone who knows the subject can't figure it out.

I have no problem with the hard-science approach. But that should be only one part of it. We also need an overview that gives the average reader a good understanding of what it is (the article on internal energy does a decent job at that). Then follow it up with the hard science.

The overview in this one borders on the estoeric and really doesn't help much.

I would be happy to re-write it... but first I have to re-discover just what the hell enthalpy is. Izuko (talk) 12:19, 6 May 2008 (UTC)[reply]

I agree that the article is more inscrutable than necessary. At least it should have a more legible lead that, instead of bombarding the reader with jargon such as "thermodynamic potential" from the first sentence, tries to put the better-known and simpler aspects of enthalpy first. For example, the definition given in IUPAC's glossary is more readable: "Internal energy of a system plus the product of pressure and volume. Its change in a system is equal to the heat brought to the system at constant pressure."[1] For many people, "heat at constant pressure" is all they need to know about enthalpy, and yet this is hidden in various places throughout the article but not mentioned in the lead. --Itub (talk) 13:19, 6 May 2008 (UTC)[reply]

I have to agree with Itub. It is likely that most readers of this article will be students who have come across the term "enthalpy" in a chemistry class. The article is written by physicists for, it seems, physicists. It is important that the chemistry reader comes across the sentence from IUPAC, "Internal energy of a system plus the product of pressure and volume. Its change in a system is equal to the heat brought to the system at constant pressure.". Can we add that to the lead? --Bduke (talk) 23:31, 20 May 2008 (UTC)[reply]

Actually despite the jargon, this article contained a lot of mistakes. I corrected most of them. Removing the jargon would be a good idea. I really doubt whether the article was writen by real experts, given the stupid mistakes it contained. I've just rewritten Helmholtz free energy. That article contained so many false statements, it was really abominable. Of course, the article now looks much more technical, but then there is really no shortcut to explaining the things from fundamental concepts. Count Iblis (talk) 23:59, 20 May 2008 (UTC)[reply]

I am aware of the errors you fixed and I have no problem with that. We should not make errors. Your comment "there is really no shortcut to explaining the things from fundamental concepts" seems to me to be a typical physicist response. There is a way. It starts from recognising who our readers are and where they are coming from. As I said, I think a lot of our readers are chemistry students. Some might well be people with even less background than that. Just think where people are going to meet the term "Enthalpy" and come to WP to find out about it. It is more than likely to be something chemical, such as heats of reaction. We should start from a simple lead, even if it is skating close to oversimplification, in order to lead people into the content. Perhaps the article should be reorganised so the chemistry student knows when to stop leaving the rest for those who want the physicists detail and absolute rigor. I have spent a career trying to teach physical chemistry to students who every year seemed to have less background in physics and mathematics than those the year before. I am not going to wade in and edit the article unless there is some kind of consensus. I spent a lot of time on Entropy maybe two years ago. It just kept reverting back to being quite unintelligible to people with a poor background in physics and mathematics. This is a major problem in dozens of articles on topics that chemistry students meet in physical chemistry courses. --Bduke (talk) 01:14, 21 May 2008 (UTC)[reply]

I do agree with rewriting this article. Thing is that the mathematical formalism and/or jargon is often just a shortcut to a correct explanation. But that correct explanation must be given. If we look at what I corrected, it was the flawed inequalities. But the reason why the mistake was made points to a fundamental misunderstanding about the basic physics. The physics that is behind this can be explained to lay persons as well using only simple examples. Take e.g. the free expansion of a gas. One can explain that as follows:

There the entropy increased despite the fact that no heat is supplied. But the gas does no work either. So, pressure times volume change (assume moving the piston extremely fast by a small amount and then lock it in place so that the gas crashes into the piston at the new locked position). One can then heuristically reason as follows: The P dv term led to the gas being accelerated, increasing the kinetic energy of the gas as it shoots out toward the piston. The gas could have performed work, but all that energy will go to waste as thermal energy in the gas.

For an ideal gas, then end result is that the temperature does not change. If you imagine measuring the temperature during the free expansion, you would actually at first see that the temperature went down as just like in case of adiabatic expansion. But then all that kinetic energy get's dissipated into heat and then the temperature get's back to where it was. That is the cause of the entropy increase you now get that you wouldn't have gotten had the gas performed work.

One can then go a step further and point out that if you make the process more violent the concept of temperature would break down. One can argue that the velocity distribution is not like that of a gas at any temperature so one needs more variables to describe the macroscopic state of the gas as it expands than just pressure, volume and temperature. But you can apply thermodynamics to the initial and final states when things have settled down. From this perspective the equations of thermodynamics becomes much more interesting because it allows you to compute the outcome of complicated processes.

The problem is that the conventions used in wikipedia is to explain some subject comprehensively in a few pages. Then you cannot spend a page to discuss the free expansion experiment in detail to make some point about thermodynamics. But but perhaps we need to change this and do what is necessary to let people really understand the topic. Count Iblis (talk) 02:24, 21 May 2008 (UTC)[reply]

Unfortunately, it seems the article has been written and re-written by laymen or by chemistry students with a very poor knowledge/ understanding of Thermodynamics and the mathematics that go with Thermodynamics. I'll try and fix it when I have the time, but in the meantime, if any professors of thermodynamics wish to edit this article, please feel free. JC 86.178.174.199 (talk) 23:08, 27 May 2014 (UTC)[reply]

Or even worse still, written by biologists!!! 86.178.174.199 (talk) 22:58, 28 May 2014 (UTC)[reply]

This article containes terrible mistakes :(

See the discusion here. You can also check my rewrite of the derivations in Helmholtz free energy and of Fundamental thermodynamic relation. I'll rewrite the relevant derivations later in the week if no one else has the time to correct them. It should be understood that keeping these errors in the page does enormous damage to the reputation of wikipedia. Count Iblis (talk) 02:25, 19 May 2008 (UTC)[reply]

I have corrected the mistakes. Count Iblis (talk) 19:53, 20 May 2008 (UTC)[reply]

While I neither bother, nor presume, to check your remediation, please accept this (however belated) verbal and virtual barnstar.
--JerzyA (talk) 16:21, 20 June 2019 (UTC)[reply]

Constant Pressure in the Definition

The "original definition" section has a question "?? Except that the pressure in a combustion cylinder is far from constant! Maybe the example is referring to a steam engine with 100% duty cycle?"

That is true, the pressure in a practical cylinder is far from constant if the internal reaction is a typical explosive combustion. If one wished to determine the work done by the gas during combustion by considering the pressure on the interior face of the piston, one would certainly have to know the detailed history of pressure vs. distance; a constant pressure could not usually be assumed. However, it is simpler to observe that a long time after the combustion is complete, when the system has returned to a quiescent state, the work done by the gas on the interior face of the piston must equal the work received by the environment at the exterior face. If the exterior face is in contact with a very large reservoir, e.g. the Earth's atmosphere, and the combustion isn't too fast, the pressure there is so close to constant that constant may be safely assumed.

Of course, a practical cylinder has a massy piston, and if the combustion is explosive, the piston may move quickly enough for the gas on its exterior face to be out of equilibrium with its surroundings. There may be sudden or oscillatory motion, producing either sound or shock waves in the environment subsequent to combustion. This real occurrence is inconsistent with assuming a constant exterior pressure. Nonetheless, for the purposes of a definition, the custom is to assume an ideal cylinder with a massless piston, and ideal combustion in which heat is produced slowly enough to maintain equilibrium of the exterior gas. The difference between real and ideal situations must be taken into account for experimental work, but is conventionally neglected for theoretical work.

The difference in sign between work done and work received was neglected in the text. This resulted from neglecting: a complete definition of work dW = F dot-product dX; whether the force under discussion is that exerted on the piston by the exterior gas or on the exterior gas by the piston; and a definition of the direction of increasing x. The positive signum shown in the text applies if one chooses these definitions: the force is that exerted by the piston on the exterior gas, and x increases on the exterior side away from the piston.

128.149.22.239 (talk) 18:55, 8 September 2009 (UTC)Jay Breidenthal jbreid@integrity.com[reply]

Fluid

This entry's basic definition uses the word "Fluid" but i believe enthalpy is not specific to fluids;just a system. Fluids are simply commonly used when measuring enthalpy. Could someone weigh in on this? - tnx JScribner (talk) 22:40, 14 September 2009 (UTC)[reply]

I agree. Internal energy is defined in terms of a thermodynamic system. Enthalpy should be defined the same way, and thermodynamic system is more accurate than fluids. I will make the change. Dolphin51 (talk) 01:45, 15 September 2009 (UTC)[reply]

Standard condidtions

I think the link Standard conditions redirects to Standard conditions for temperature and pressure. From what I understand enthalpy is calculated at a constant pressure not pressure and temperature. Shouldn't it redirect to standard state? --kupirijo (talk) 10:36, 25 October 2009 (UTC)[reply]

Changed a sentence

There was a sentence that said "the change in enthalpy at constant pressure is equal to the heat added plus the work done". That is false. I changed it to "the change in enthalpy at constant pressure is equal to the heat added".

After all, that is why the concept of enthalpy even exists, it was because for the scientists in previous centuries with bad equipment, it let them run their experiments with pistons at constant pressure. Enthalpy eliminated a variable "work". This made life easier for them, so they created the concept of enthalpy. —Preceding unsigned comment added by Apc3161 (talk • contribs) 17:10, 5 January 2010 (UTC)[reply]

Sudden change of notation

in the 'Difference between H and U: the additional term pV' the writer suddenly introduces a variable uppercase P to represent pressure [as far as i can tell]. I'm almost certain that lowercase 'p' is intended, but the switch is confusing. Perhaps someone with a more complete knowledge of the subject would have the confidence to make the appropiate modifications. Hai2410 (talk) —Preceding undated comment added 01:00, 24 January 2010 (UTC).[reply]

It's always been P in my experience. JIMp_talk·cont 22:55, 6 June 2011 (UTC)[reply]

Non-mechanical = heat??

In the beginning of this wiki, the term "non-mechanical" is used 6 times. So, I assume this means heat energy or something kind of like the type of energy when a photon of energy enters an atom system? Does this mean heat is going into the system, or out of? Clarification would be helpful. —Preceding unsigned comment added by 71.193.147.120 (talk) 20:40, 7 June 2010 (UTC)[reply]

The statement "ΔH of a system is equal to the sum of non-mechanical work done on it and the heat supplied to it." is a bit baffling. It is not clear what the writer meant by "non-mechanical work". From the definition of H = U + PV, it follows that dH = dU + PdV + VdP. The heat flow is δQ = dU + δW where δW is the work done BY the system. For reversible processes, δW = PdV so dH = (dU + PdV) + VdP = δQrev + VdP. This means that ΔH = Qrev + ∫VdP. So, ΔH is equal to the sum of the integral of VdP and the reversible heat flow. So it seems to me that the statement as it stands is wrong, as well as being unclear. AMSask (talk) 00:50, 4 June 2014 (UTC)[reply]

External fields

This version of the article says, "In the absence of an external field, the enthalpy may be defined, as it is generally known, by: H = U + pV". I have removed the bit about the absence of the external field because it seems the definition should still hold even in the presence of an external field. Either the external field affects the internal energy, pressure, and/or volume; or it does not. In either case I cannot envision a situation where the presence of an external field would make this definition incorrect. Anyone else want to weigh in on this? Riick (talk) 19:34, 16 August 2010 (UTC)[reply]

Enthalpy is Not Heat

There's a contradiction here, and I don't have the grounding in this subject to be able to fix it.

   Adding or removing energy through heat is the only way to change the enthalpy, and
   The amount of change in enthalpy is equal to the amount of energy added through heat.

... However, heat is not the only way to change enthalpy.

Maybe I'm missing something? As I'm re-reading, I'm getting that my interpretation of the section is incorrect, but I can't figure out why. Could use clarification. — Preceding unsigned comment added by 108.57.35.49 (talk) 15:27, 31 May 2011 (UTC)[reply]

No, enthalpy H is not heat, H = U + PV. 86.168.55.23 (talk) 02:08, 20 October 2012 (UTC)[reply]

H is a state function, q is a path function, however at constant pressure (and where only expansion/contraction work is allowed) the mathematics show that dq = dH, therefore dq is proved to be a state function under the special case of constant pressure, whereas generally speaking q is a path function, and also in this special case change in heat (dq) is equal to the change in enthalpy (dH). 86.168.55.23 (talk) 02:13, 20 October 2012 (UTC)[reply]

While I shan't waste time critiquing the pedagogy head-on, failing to link path function, enthalpy (and state function) is at least supercilious, even if they're already linked in the article (tho I also won't further decorate this talk page, beyond the arguably redundant remark that the connecting algebra raises this question: was the deeper purpose of that wording to either mock or preen? ... oh, well: for that matter, "Look at the big vocab on Jerzy".)
~~----
— Preceding unsigned comment added by JerzyA (talk • contribs) 15:59, 20 June 2019 (UTC)[reply]

Changes by Adwaele on June 21, 2012

In the Section “Formal definition” it is stated that p is the pressure at the boundary of the system and its environment. This is wrong/unclear for the following reasons:

The pressure need not be the same all over the boundary of the system. E.g. in the case of throttling usually an open system is chosen that contains the plug. In this case the pressures at the inlet and outlet differ. Which pressure should be taken in this case?
This formulation leaves unclear which pressure (and temperature?) must be chosen in calculating U. Is this also the value at the boundary?
In an inhomogeneous system the value of the enthalpy of the total system depends on the pressure and temperature distribution inside the system and not only on the pressure at the boundary.

The Article refers to Zumdahl, but Zumdahl talks about “the” pressure of the system (suggesting a homogeneous value) and not about the pressure at the boundary of the system and its environment.

The solution is simple: first define the enthalpy for homogeneous systems. This has also the advantage that δQ=TdS, since no irreversible processes can take place in a homogeneous system. Also the pressure is well-defined so δW=pdV. Next use the property that enthalpy is an extensive parameter. So the enthalpy of an inhomogeneous system is the sum of the enthalpies of the subsystems. In case of continuous variation of p and T composition etc. the summation becomes an integration.

The Subsection “Difference between enthalpy and internal energy” is not part of the formal definition and should be moved to a separate Section.

I have modified the text accordingly.

I like to add that there is a lot of overlap between the various Sections that needs to be cleaned up.

(Adwaele (talk) 21:05, 21 July 2012 (UTC)).[reply]

The article is wrong

This article is wrong starting from the beginning: "Enthalpy is a measure of the total energy of a thermodynamic system." Enthalpy is not a measure of the total energy of a thermodynamic system. It is an artificial mathematical function H where H = U + PV, where U is the total energy of a thermodynamic system, and P is the pressure and V is the volume. The beauty of this artificial mathematical function is that thermodynamic experiments are usually carried out to measure changes in a quantity against one variable whilst holding other variables constant, and useful properties of the function themselves can be proved through the mathematics. For example, dq can be shown to be a state function and not a path function as might be expected from the property of q, under constant pressure. JC86.168.55.23 (talk) 02:04, 20 October 2012 (UTC)[reply]

Agreed: The first sentence of the present article "Enthalpy is a measure of the total energy of a thermodynamic system." is not correct. This should be changed. I can't believe that this error has remained for so long. No cite is provided. Where did this statement come from?AMSask (talk) 13:52, 1 April 2014 (UTC)[reply]

I would suggest a change to the first paragraph:

Enthalpy is a defined thermodynamic potential of a thermodynamic system consisting of the internal energy of that system plus the product of pressure and volume of the system:

H=U+PV

The unit of measurement for enthalpy in the International System of Units (SI) is the joule, but other historical, conventional units are still in use, such as the British thermal unit and the calorie:

Since enthalpy, H, consists of internal energy, U, plus the product of pressure, P and the volume of the system, V, which are all functions of the state of the thermodynamic system, enthalpy is a state function.

While the units of enthalpy are units of energy, the enthalpy of a system should not be confused with internal energy, which is the total energy of a system.AMSask (talk) 19:23, 1 April 2014 (UTC)[reply]

The historical significance of putting H = U + PV by definition as an artificial mathematical function is as follows: During the Industrial Revolution, engines powered by burning coal to do work were invented and built. People were building better engines over time, better in the sense that more efficient engines, ie more work for the same amount of coal consumed, were built. So people wanted to know whether there was a limit on how much work you can get out from burning a fixed amount of coal. Obviously you cannot prove this one way or another by building more and more engines, because say in the first year, a train engine moved 1 mile by burning 1 tonnes of coal; in the second year a more efficient engine was built and it gave a mileage of 2; in the third year a better engine was built giving a mileage of 3, etc. So using this projection an engine built a thousand years later should give a mileage of 1000. What was needed was a mathematical way of making a proof one way or the other. By setting H = U + PV (a state function), it was proved that dq = dH at constant pressure. Therefore it was proved that dq was in fact a state function. That is to say, the heat change by a system moving from one state to another state is a fixed amount. ie Only a fixed amount of heat (which can only be turned into a fixed amount of work) can be obtained by burning a tonne of coal. JC 86.178.174.199 (talk) 00:41, 27 May 2014 (UTC)[reply]

Someone has made the opening even worse: "Enthalpy is a defined thermodynamic potential". Enthalpy is definitely not a defined thermodynamic potential. Enthalpy H is simply defined as U + PV. It may look simple, but H is a function of three variables, U, P and V, so is mathematically sophisticated. An example of a thermodynamic potential would be dH/dT (where T is temperature) or dH/dM (where M is mole quantity), (derivatives and partial derivatives). Please would only people who understand Thermodynamics and its associated mathematics edit the article. I am going to change this opening sentence. JC.86.178.174.199 (talk) 00:55, 27 May 2014 (UTC)[reply]

JC: I am not sure what your statement "Enthalpy is definitely not a defined thermodynamic potential." is based on. The 4 functions U, H, G, F are commonly referred to as thermodynamic potentials. See, for example: http://wiki.riteme.site/wiki/Thermodynamic_square AND http://hyperphysics.phy-astr.gsu.edu/hbase/thermo/thepot.html#c1 AMSask (talk) 01:17, 31 May 2014 (UTC)[reply]

Nope. Enthalpy is not thermodynamic potential. It is a thermodynamic state function. A potential in physics is the energy change per unit change of something else, and its units are joules per something, eg, joules per metre, joules per coulomb, joules per degree, joules per mole, etc. The unit of enthalpy is joules. AMSask may have confused enthalpy with the chemical potential, mu, (I apologise but I don't have Greek fonts), and the chemical potential (units of joules per mole) = the following (again I apologise for not having the partial derivative font, so I'll just use d with conditions being kept constant explicitly stated, where n indicates a given substance:

chemical potential for substance 1, mu(1) = (dG/dn1)at constant pressure, temperature, all other substances eg n2, n3 etc

                                          = (dU/dn1) at constant volume, entropy, n2,n3,n....

                                          = (dH/dn1) at constant pressure, entropy, n2, n3, n....

                                          = (dA/dn1) at constant volume, temperature, n2, n3, n....

As you can see one of the definition of chemical potential is dH/dn with conditions.

Ref: Peter Atkins, Physical Chemistry second edition reprinted 1984, ISBN 0-19-855150-9, pages 175 &176

109.158.28.28 (talk) 00:19, 17 November 2015 (UTC) JC[reply]

Just to point out to AMSask, the article in Wiki he/she cites titled thermodynamic potential does not cite any reference on where the title came from. 109.158.28.28 (talk) 00:26, 17 November 2015 (UTC)[reply]

I've changed the opening 2 paragraphs, only to discover the remainder of the opening is also gibberish. I'll try and fix it when I have the time. The Laws of Thermodynamics are general and not just for chemical reactions or changes. JC86.178.174.199 (talk) 01:14, 27 May 2014 (UTC)[reply]

JC: You should use your Wikipedia Login so we know who you are. This article has had a lot of problems in the past and we would welcome any thoughtful corrections you might wish to make. But given the history of this page it would be best to provide authorities for any changes. You are making changes without citing any authorities and some of your changes appear to be wrong. Your definition of "thermodynamic potential" appears to be inconsistent with common usage. Please provide a cite to an authority for your statements. AMSask (talk) 13:20, 30 May 2014 (UTC)[reply]

AMSask, are you actually a Physical Chemist? Do you know any thing about Thermodynamics? H and G are mathematical functions (I don't know the F to which you refer); they are sometimes described as chemical potentials when dealing with chemical thermodynamics. However, Thermodymamics theory is general and is not simply restricted to chemical systems. Many concepts or descriptions in physics use explanations that are incorrect, for example the idea of "flux" in electromagnetism. These explanations are useful so long as you know what they are trying to depict, and making the maths easier to perform. If this article in Wiki cannot be trusted, how can you trust other wiki pages on the subject? 86.178.173.142 (talk) 01:50, 1 June 2014 (UTC)[reply]

JC. Please identify yourself with a Wikipedia Login. I am not sure what requiring citations to authorities has to do with one's qualifications. If something is correct you should be able to provide a cite for it. Your statement that H is not a thermodynamic potential requires a cite. It is incorrect. I have provided you with links to two pages so you will understand what is meant by the term "thermodynamic potential". You seem to be using a different concept.AMSask (talk) 12:13, 1 June 2014 (UTC)[reply]

AMSask, since you asked, qualifications are very important in a technical subject. It is easy to see that you do not understand what you read, let alone being able to assess what you read. Please do us a favour and not continue to confuse students or biologists who read this article. 109.158.28.28 (talk) 00:42, 17 November 2015 (UTC)[reply]

F is the Helmholtz free energy, also sometimes designated as A: F = U-TS. See: https://wiki.riteme.site/wiki/Helmholtz_free_energy AMSask (talk) 12:13, 1 June 2014 (UTC)[reply]

Yes, we use A over here to save confusion as F is also used to designate a function. 86.178.173.142 (talk) 23:21, 1 June 2014 (UTC)[reply]

Your statement: "Enthalpy change accounts for heat transferred to (or from) the environment at constant pressure, with measurements data quoted at constant temperature. requires a citation. The first part (non-bold) is correct. Can you provide an explanation for the bold part and a cite for the bold part? If not, it will be removed.AMSask (talk) 12:21, 1 June 2014 (UTC)[reply]

This is known by everyone in the business. You can of course measure the initial state at a constant temperature T1, and then the final state at another constant temperature T2. 86.178.173.142 (talk) 23:21, 1 June 2014 (UTC)[reply]

"Constant temperature" means the temperature does not change during a process ie. constant over time while the process is occurring. It is very confusing to students if you try to use "constant temperature" to mean "uniform temperature" ie. over volume. It is implicit that the system has a uniform temperature throughout its volume because that is how temperature is defined for a thermodynamic state (equilibrium). It is redundant if you are taking it to mean the temperature is constant while in a constant thermodynamic state. Please remove this or clarify and provide a mainstream authority that uses "constant temperature" in this way. AMSask (talk) 13:54, 2 June 2014 (UTC)[reply]

JC: Unless you can provide a cite for your statement that "Enthalpy is an artificial mathematical function" it appears to be simply your POV. The cite for "Enthalpy is a defined thermodynamic potential" is contained in the link to the Wikipedia page on "thermodynamic potential". AMSask (talk) 12:25, 1 June 2014 (UTC)[reply]

AMSask H = U + PV is invented by a mathematician. It is a piece of creative accounting, but extremely useful and beautiful. It is not my POV, it is the POV of everyone in this business. 86.178.173.142 (talk) 23:21, 1 June 2014 (UTC)[reply]

Please provide your reference for the statement that 'a mathematician' invented H. The objection to the term "artificial mathematical function" is that it is non-standard usage, that it is ambiguous (what do you mean by "artificial"?) and it is confusing. "Artificial" suggests something that is not natural. Can you provide a reference to a text or peer reviewed material that uses it in this way? If not, it seems to violate Wikipedia NPOV policy. Back to my question that you seem to be avoiding, how can you maintain that H is not a thermodynamic potential? See: Wikipedia article on Thermodynamic potential and the references cited. This is a well defined and well understood concept in thermodynamics. AMSask (talk) 13:54, 2 June 2014 (UTC)[reply]

JC please use a Wikipedia Login ID. Your IP address keeps changing and it is difficult to know whether you are the same or different person making the comments. AMSask (talk) 13:54, 2 June 2014 (UTC)[reply]

I suggest AMSask get himself a copy of Peter Atkins "Physical Chemistry"- various editions, and read and do the problems set. 86.135.49.164 (talk) 00:21, 9 October 2015 (UTC)[reply]

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Suggestion on how to make concept accessible to non-technical readers

My first glimmer of understanding was when I saw "BTU." I moved the units to the first paragraph.

I came away with a still-tenuous grasp of the concept. My learning mind needs a concrete example to then be able to grasp a generalization A single concrete example at the beginning could go a long way to giving a basic grasp of the concept to non-technical people who don't already have it.

After a quick peruse, the below candidate explanatory example is my best guess of what an example intro could be, but I defer to someone who understands the concept better than me to provide a useful example.

For example, a canister of compressed butane under pressure at room temperature can be thought of having two kinds of energy. One is the potential energy from release of the pressure of the gas inside the canister. This can do work such as propelling gas out of the canister until pressure is equalized with the room. The second is the latent chemical energy of the butane, which is released when the butane is burned. A rough measure of the enthalpy of the system can be calculated by measuring the units of energy represented by the pressure and chemical energy in the butane canister and adding them together.

This example reflects that after a quick read, I wasn't able to exclude chemical energy as part of the measurement of the enthalpy of the system. If it's not, an example that says enthalpy does not include the chemical energy from burning the butane would be very helpful, since I think of calories and BTUs in terms of how much a given quantity of water can be heated by burning a specific quantity of gas.

Including a very basic correct(ed) example like a charged butane cannister may seem like stating the obvious, but that is exactly what is needed by readers who come to the article without any clue of what enthalpy is, and absent an example leave in the same state. Which may be a rather large percentage of the visitor to the page.

I suggest a concise very concrete example, very early in the article. - Preceding edit added by User talk:Ocdcntx on 18 February 2018.

Enthalpy is not energy so an explanation based on adding different kinds of energy won’t be the example you are seeking. (It might be helpful to observe that momentum is also not energy, but momentum is a very useful property nonetheless.)

I can provide a sound explanatory example although I would hesitate before inserting it in Wikipedia because Wikipedia isn’t a text book. If J joules of heat flow into a solid or liquid, the internal energy of the solid or liquid increases by J joules. Conversely, if the internal energy of a solid or liquid increases by J joules we can say with confidence that J joules of heat have flowed into it.

Similarly, if J joules of heat flow into a body of gas constrained at constant volume, the internal energy will increase by J joules.

Now imagine that the constraint is removed from the body of gas, allowing it to expand so that its pressure returns to its original value. As it expands it does N newton.metres of positive work on the surroundings. Also its temperature is reduced, but not to its original temperature. In accordance with the first law of thermodynamics the overall change in the internal energy of the gas is J minus N (the heat added, minus the work done.)

The internal energy, plus an allowance for the work that can be done, is called enthalpy. So if the enthalpy of a body of gas increases by J joules as it is heated at constant pressure, we can say with confidence that J joules of heat flowed into the gas. Conversely, if J joules of heat flow into a gas at constant pressure, the increase in enthalpy will be J joules even though the increase in internal energy will be less than J joules.

The “allowance for the work that can be done” is quantified by PV, the product of the pressure and the volume of the body of gas. Enthalpy H is equal to internal energy U plus PV:

H = U + PV

Dolphin (t) 23:27, 18 February 2018 (UTC)[reply]

Nonstandard term matter transfer/flow

I have seen that the article is littered with the nonstandard term matter transfer/flow claimed by some editor (only in an edit description!) to be different than ordinary/usual mass transfer. The claimed difference (if any) is not clearly explained in article if it is claimed to be an important distinction. I have removed some occurrences of the nonstandard term which should not be used in a wikiarticle. I may have missed other cases of use of this nonstandard language. I'll remove them when spotted.--37.251.221.101 (talk) 23:26, 17 August 2019 (UTC)[reply]

A small survey (alphabetically stopping at 'G': Adkins 3rd edition, Buchdahl, Callen 2nd edition, Denbigh 4th edition, Guggenheim 5th edition; for now I will not burden this page with the extensive quotes to make this a full account) of standard reliable sources on thermodynamics as a branch of physics finds no instances of the term 'mass transfer'. Terms such as 'transfer of matter' and 'transfer of material' are found instead. There are abundant examples of the shift from 'transfer of xxx' to 'xxx transfer', where 'xxx' is used as a noun in the former and as an adjective in the latter. Callen is a reliable source for such questions; he uses the term 'matter transfer' twice. True, the term 'mass transfer' is the title of a Wikipedia article on that topic, but that does not make the term standard in thermodynamics. It is not our task to invent standardization of terms. A Wikipedia article is not a reliable source.Chjoaygame (talk) 19:55, 18 August 2019 (UTC)[reply]

I have seen the mentioned survey of expressions from sources containing matter transfer/transfer of matter (from an editor sandbox). The mentioned wordings in sources are an example of sloppy use (in ordinary language) of the word matter to mean substance or material. Such sloppy/careless uses of words in ordinary language are not acceptable when scientific terminology is involved, ordinary language is very frequently affected by ambiguities and use of words having the meaning of other words. Scientific language is a rather formal(ized) language or controlled language because the inconsistencies and ambiguities of natural ordinary language need to be purged. Similarly, some editor's fascination (or obsession?) with ordinary language and its weak features is also not acceptable for scientific articles on Wikipedia.--37.251.220.113 (talk) 13:26, 20 August 2019 (UTC)[reply]

Another aspect to be analyzed involves the claimed lack of standard status of the term mass transfer in thermodynamics, despite its use in transport phenomena. It is claimed that the term mass transfer in the usual sense re transport phenomena is not accepted in thermodynamics as the sloppy (and undefined) expression transfer of matter is considered somehow distinct in sense from mass transfer. I don't understand the supposed distinction in terms between the transport phenomena and thermodynamics, as if the non-equilibrium thermodynamics which includes transport phenomena were non-existent.--37.251.220.113 (talk) 13:58, 20 August 2019 (UTC)[reply]

variable of state

There are several ways of considering enthalpy.

It is said to be a function of state, but also to be a state variable.

If the word 'function' here is to have its usual mathematical meaning, then enthalpy as a function of state would have arguments, with enthalpy evaluated for a specific set of values of the arguments. This way of talking is not much in evidence in the article.

In more detail, the article writes;

H=T\,dS+V\,dp+\sum _{i}\mu _{i}\,dN_{i},\,\,\,\,\,\,\,(1)

This naturally leads to the form of expression of $H$ as a mathematical function

H(S,p,\{N_{i}\}),\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,(2)

where $S,p,$ and $\{{N_{i}}\}$ denote the obvious variables of state, called the natural arguments of the enthalpy, the arguments of that function of state.

That form of expression is a Legendre transform of

U(S,V,\{N_{i}\}),\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,(3)

which exhibits the internal energy as another function of state, with different but related arguments, also state variables.

The two arguments denoted by $S$ in the two formulas (2) and (3) are measured under different conditions, and more properly would have different symbols, but I will pass over that for brevity here.

For reasons that need not be stated here, knowing the form (2), one can also write

S(H,p,\{N_{i}\}),\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,(4)

which exhibits $S$ as a function of state variables $H,p,$ and $\{{N_{i}}\}$ .Chjoaygame (talk) 07:42, 5 February 2020 (UTC)[reply]

? Comment about Gibbs free energy?Chjoaygame (talk) 21:31, 5 February 2020 (UTC)[reply]

Callen gives a systematic account of the energy functions such as internal energy, enthalpy, Helmholtz free energy, Gibbs free energy. There are dozens of such functions, one for each combination of controlled state variables. As controlled state variables, internal energy uses entropy, volume, and suitably specified chemical composition. Enthalpy uses entropy, pressure, and suitably specified chemical composition. Processes can be classified by the state variables that are to be varied or kept constant. For a process of adding energy purely as heat, at constant suitably specified chemical composition, at constant volume, internal energy is convenient; in such a process, the system is prevented from gaining or losing energy as thermodynamic work during the heating, while its temperature changes. For a process of adding energy as heat, at constant suitably specified chemical composition, at constant pressure, allowing also the system to gain or lose energy as thermodynamic work by changing its volume, enthalpy is convenient; in such a process, in general, but depending on its chemical composition, the system temperature will also change.Chjoaygame (talk) 11:27, 12 July 2020 (UTC)[reply]

undone edit

I undid this edit. Its edit note said "Deleted the false statement that 'a system enclosed so as to prevent mass transfer' is a requirement for the heat transfer at constant pressure to be equal to the change in enthalpy. A reaction occurring in an open container at atmospheric pressure would also undergo heat transfer equal to the change in enthalpy."

The sentence in the article, that the undone edit note says was false, did not say that a closed system is a requirement. It said that when a system is closed then such and such holds, but did not say that closure is a necessary condition.

For identification of heat transfer in a vessel that has a path that allows matter transfer, the heat transfer needs to be by separate path that does not allow transfer of matter.Chjoaygame (talk) 00:45, 1 May 2020 (UTC)[reply]

Clarification on the opening definition

When every introductory chemistry textbook gives the example of deltaH = q at constant pressure as a reaction occurring in an open container, would it not be prudent to state why that it should be enclosed?. Whilst it may be a tacit assumption for those with wider knowledge in the area, its nitpicky and confusing from a pedagogical point of view. — Preceding unsigned comment added by 2A00:23C6:B106:4600:8906:64C8:7E10:2A16 (talk) 15:51, 4 May 2020 (UTC)[reply]

I made a change that I hope may be a reasonable compromise, from "enclosed" to 'contained'.Chjoaygame (talk) 16:54, 4 May 2020 (UTC)[reply]

Notation of molar and specific enthalpy

In my thermodynamics textbook (Koretsky Thermodynamics), Enthalpy has the symbol H, specific enthalpy has the symbol ĥ and molar enthalpy has the symbol h. I notice that in this Wikipedia page it's defined as H for enthalpy, H_m for molar enthalpy, and h for specific enthalpy. Can I ask where this notation comes from? — Preceding unsigned comment added by Jason0597 (talk • contribs) 16:48, 20 May 2020 (UTC)[reply]

Different authors use different notation, so in a Wikipedia article with many sources we have to choose one notation for each quantity and stick with that even if some equations don't look the same as in the sources they came from.

I have not seen Koretsky's book. However in the books I know H_m for molar is fairly standard but the symbol for specific varies widely. In Atkins' Physical Chemistry (8th edition) for example, molar enthalpy is H_m but specific enthalpy is actually given no symbol; he mentions values for a few combustions and either calls them ΔH_m/M or else writes out "specific enthalpy of combustion". Dirac66 (talk) 19:24, 24 May 2020 (UTC)[reply]

Talk of edit and undo

I undid this edit that was posted by Editor Koavf. It was good faith but inappropriate.

The undone edit replaced an entry that served to point out the need to distinguish enthalpy from entropy. The former entry read "Not to be confused with Entropy." It is appropriate to point that out in that place at the head of the article, but not to try to explicate the nature of entropy there. Such an explication belongs elsewhere.

The entry that I undid read "For the tendency toward uniformity and the breakdown of information in the universe, see Entropy." Entropy is a physical quantity, not a tendency. The tendency is asserted by the Second Law of Thermodynamics.Chjoaygame (talk) 15:21, 10 July 2020 (UTC)[reply]

Chjoaygame, I recommend making a better explanation than removing it. {{distinguish}} doesn't explain what the problem is, whereas {{for}} does. ―Justin (koavf)❤T☮C☺M☯ 09:12, 11 July 2020 (UTC)[reply]

I think that the article on enthalpy is not the place to try in a few words to explain what entropy is. I think that all that is suitable for that place is a distinction with a link to a better explanation.Chjoaygame (talk) 09:20, 11 July 2020 (UTC)[reply]

Chjoaygame, Then use the same short description as on entropy. You're forcing users to read the two articles to understand if they're on the correct page. ―Justin (koavf)❤T☮C☺M☯ 10:15, 11 July 2020 (UTC)[reply]

”tendency toward uniformity” is one brief summary of entropy; and “breakdown of information” is another, but these are not the only two, and they are esoteric explanations. Any readers who are uncertain about whether they should be reading Enthalpy or Entropy are unlikely to find assistance from these two, esoteric, explanations.

‘’A thermodynamic property’’ is the one that means the most to me but I’m not suggesting we add this one in an attempt to cover all known meanings of entropy. I’m inclined to agree with Chjoaygame and say “not to be confused with Entropy.” Dolphin (t) 14:27, 11 July 2020 (UTC)[reply]

Done.Chjoaygame (talk) 22:08, 11 July 2020 (UTC)[reply]

Such pointers of distinction between enthalpy and entropy at the top of the article are completely useless. Why would someone confuse the two? Not by misspellings, for example, and they do they rhyme? If the reader doesn't know what they are looking for, then neither term by itself will clarify anything. Placing long explanations of the term's relation to other coordinates is an lengthy game, the list could be long before they recognize what they want. So, instead of tempting them to click around the web or Wikipedia, let them read the article's lead, that should clarify the subject matter without ambiguity and without excursions to other places. Kbrose (talk) 18:44, 12 July 2020 (UTC)[reply]

I think the "Not to be confused with entropy" may be useful to readers who have never studied thermodynamics, even at the level of general chemistry. I agree that there is no point in trying to give a one-line summary of a concept as complicated as entropy. But at least for the beginner who looks here for an explanation of entropy(!), we can briefly tell him or her that enthalpy is not the same as entropy, and also provide a link to the correct article. Dirac66 (talk) 19:15, 12 July 2020 (UTC)[reply]

It seems to me that in this case, such a tag might confuse a novice even more, because it implies that it is indeed a common mistake to confuse the terms, when it is not. To get an overview of related topics in the field the article has a sidebar for thermodynamic, where more context is provided by the groupings. Kbrose (talk) 21:24, 12 July 2020 (UTC)[reply]

Having taught physical chemistry in universities for 40 years in 4 different countries, I can assure you that indeed many students confuse enthalpy and entropy. --Bduke (talk) 22:54, 12 July 2020 (UTC)[reply]

@User:Bduke. I don’t think anyone would challenge the idea that students can confuse the two. The question before us is which is the best way to guide students to the article they are looking for – encourage them to peruse the lead, or to use a hatnote saying entropy is a tendency toward uniformity, and breakdown of information. My view is that the hatnote won’t be of much use to readers whose knowledge is so basic that they confuse the two. Dolphin (t) 02:46, 13 July 2020 (UTC)[reply]

Thank you, I agree that this formulation expresses my view more fully as well. I also think that it is not Wikipedia's task to guide students. Wikipedia is not intended to be a text book. And therefore Wikipedia should not have statements at the top of articles that reminds novice readers that they don't know what they are doing. Sometimes there is legitimate concern about confusion, for example the difference between premise and premises comes to mind. Kbrose (talk) 16:01, 13 July 2020 (UTC)[reply]

Yes, premise and premises can be confused because their written forms are so similar. And two thermodynamic functions named ent...py and ent..py also can confuse readers. Dirac66 (talk) 17:36, 13 July 2020 (UTC)[reply]

Wikipedia isn’t a text book, but the same problem confronts the authors of text books on thermodynamics, and the students who use those text books. One chapter titled Enthalpy, another titled Entropy, and no hatnotes at the tops of chapters. Everyone copes with this situation. Dolphin (t) 00:10, 14 July 2020 (UTC)[reply]

A textbook has a table of contents short enough so that the reader who is looking for the concept will note the presence of 2 chapters with similar names. If he is not sure which is which, the table of contents will suggest that he look at both.

However English Wikipedia now has over 6 million articles so no one can peruse a full table of contents. For the poor reader who vaguely remembers hearing about a measure of disorder called something like "enthopy" (mis-spelling deliberate) and is confused that this article doesn't seem to fit, it seems kind to suggest trying Entropy. But not to try to summarize what Entropy is in one line, because the concept is just too complicated for that. Dirac66 (talk) 01:37, 14 July 2020 (UTC)[reply]

new edit of lead section

This new edit writes

For example the heat absorbed by a chemical reaction at constant pressure and temperature is equal to the enthalpy change for the reaction, and is known as the standard enthalpy of reaction.

The Gibbs free energy change is measured at controlled temperature and pressure. Is that relevant here?Chjoaygame (talk) 03:54, 26 February 2021 (UTC)[reply]

I don’t claim any familiarity with Gibbs Free Energy and the sentence quoted above makes sense to me.

The sentence mentions “constant pressure and temperature” and that expression includes both standard P and T; and controlled P and T. There may be value in adding a comment about Gibbs Free Energy, but I don’t consider it essential to the concept of standard enthalpy of reaction. Dolphin (t) 05:28, 26 February 2021 (UTC)[reply]

A reliable source for the sentence would perhaps help.Chjoaygame (talk) 21:41, 26 February 2021 (UTC)[reply]

Chjoaygame and Dirac66: The edit in question added a sentence to the lead. The lead is intended only to provide an easily understood summary of topics that appear in detail later in the article. It is acceptable, and the norm, for in-line citations not to be used in the lead.

On close inspection I am now of the view that the sentence in question does not have a place in the lead – it is too specific for the lead (which is intended to provide a general summary) and it appears to supply information that is not found elsewhere in the article. I think this sentence should be moved to a more appropriate position later in the article and it should be adequately supported by citation of a reliable published source. Dolphin (t) 00:25, 27 February 2021 (UTC)[reply]

The added information is found and sourced in the section Heat of Reaction, with only slightly different wording. The brief mention does belong in the introduction because most chemistry textbooks introduce enthalpy by relating it to heats of reaction. Dirac66 (talk) 02:13, 27 February 2021 (UTC)[reply]

The new information has been added to the third paragraph of the lead. This paragraph is correctly written to apply generally to changes in enthalpy; the new information is not of a general nature – it contrasts with the rest of the third paragraph by being very specific on two counts. Firstly the new information focuses on chemical thermodynamics whereas the rest of the paragraph is also applicable to the thermodynamics of heat engines and refrigerators and all other applications of enthalpy. Secondly the new information is stated to be applicable in the case of constant pressure and constant temperature; the restriction to constant temperature is a very specific application of enthalpy; enthalpy is generally applicable to processes in which the temperature does not remain constant.

I am in favour of some or all of the new information being moved to a more appropriate location in one of the detailed sections of the article, possibly the section titled “Heat of reaction”. Dolphin (t) 14:19, 27 February 2021 (UTC)[reply]

Since it contrasts with the rest of the third paragraph, I have now moved it to a separate paragraph of the lead. But I still think it belongs somewhere in the lead, since it is the first brief mention and summary of enthalpy in chemistry. Yes, more detail and sources are found in the section "Heat of reaction". Dirac66 (talk) 01:36, 28 February 2021 (UTC)[reply]

Thanks, your change is an improvement – both moving the sentence to its own paragraph, and omitting the suggestion that it is an example.

However, by continuing to use the words “and temperature” the sentence now implies that if a chemical reaction doesn’t take place at constant temperature the heat absorbed is not equal to the change in enthalpy; and the reaction doesn’t have an enthalpy of reaction. Both implications are incorrect so I suggest the words “and temperature” must be removed. Dolphin (t) 06:56, 28 February 2021 (UTC)[reply]

I have now reworded the sentence slightly to remove the false implications. Dirac66 (talk) 20:55, 28 February 2021 (UTC)[reply]

Thinking it over. The text in question in the lead reads

In chemistry, the heat absorbed by a chemical reaction at constant pressure and temperature is of interest and is equal to the enthalpy change for the reaction, known as the enthalpy of reaction.

This relates to the following text in the body of the article.

For an exothermic reaction at constant pressure, the system's change in enthalpy equals the energy released in the reaction, including the energy retained in the system and lost through expansion against its surroundings. In a similar manner, for an endothermic reaction, the system's change in enthalpy is equal to the energy absorbed in the reaction, including the energy lost by the system and gained from compression from its surroundings. If

Δ H

is positive, the reaction is endothermic, that is heat is absorbed by the system due to the products of the reaction having a greater enthalpy than the reactants. On the other hand, if

Δ H

is negative, the reaction is exothermic, that is the overall decrease in enthalpy is achieved by the generation of heat.^[1]

^ Laidler, Keith J.; Meiser, John H. (1982). Physical Chemistry. Benjamin/Cummings. p. 53. ISBN 978-0-8053-5682-3.

It seems to me that the enthalpy is used for 'heat of reaction' at (conveniently atmospheric) constant pressure because it is intended to measure the energy of the chemical reaction by calorimetry. The pressure for the process is controlled simply by exposing the system to atmospheric pressure. The system is allowed to choose its volume change, and P–V work, which, however, is not measured, and in any case is ignored for this process. The system is allowed to gain or lose P–V energy, by being compressed by, or by compressing its surroundings. It is intended that the energy coming into or out of chemical bonding should appear and be measured as change of temperature, without letting the system gain or lose energy as heat by transfer of heat to or from the surroundings; there is no change of entropy of the system just by heat transfer to or from the surroundings. In this sense, the temperature change in the reaction is 'uncontrolled'. This is calorimetry, performed in a calorimeter. But the initial condition of the system, before reaction, has not been specified in my just foregoing sentences. The initial condition is specified by setting the initial temperature at a convenient standard value. This does not amount to controlling the temperature during the process of reaction, the final temperature being uncontrolled.

The situation is different with Gibbs free energy, for which both temperature and pressure are controlled for the process. The system is allowed to gain or lose P–V energy by being compressed by or by compressing its surroundings. The pressure is controlled simply by exposing the system to atmospheric pressure. This time again, the system is allowed to choose its volume change, which, in contrast, is now accurately measured, and provides the information that measures the energy of reaction, as P–V work. In contrast to the enthalpy case, the energy coming into or out of chemical bonding is made to appear at thermostatically controlled unchanging temperature, letting the system choose to gain or lose energy as heat by an unmeasured transfer of heat to or from the surroundings, associated with system entropy change. In this sense, the temperature change in the reaction is 'controlled'. This is not calorimetry, because the entropy change of the system is not constrained by a thermally isolating container. Again, the initial condition of the system, before reaction, has not been specified in my just foregoing sentences. Again, the initial condition is specified by setting the initial temperature at a convenient standard value. This fits with controlling the temperature during the process of reaction.

I think the text in the lead needs to take this into account.Chjoaygame (talk) 04:46, 1 March 2021 (UTC)[reply]

Perhaps

In chemistry, starting in standard conditions, the energy absorbed by a chemical reaction at atmospheric pressure is conveniently measured by calorimetry. It has been called the 'heat of reaction', and is known as the enthalpy of reaction.

Thoughts?Chjoaygame (talk) 14:27, 1 March 2021 (UTC)[reply]

Different types of calorimeter operate in different conditions. The textbook by Atkins and de Paula (Physical Chemistry, 8th edition) considers first the constant-volume bomb calorimeter (p.38) which actually measures ΔU(combustion), to which one adds Δ(PV) to obtain the more useful ΔH. Then on p.41-42, this text mentions isobaric calorimetry which determines ΔH, including both adiabatic flame calorimetry and differential scanning calorimetry. Of course we cannot explain all this in the lead, but if we mention calorimetry I think we should agree what type(s) we are talking about. Dirac66 (talk) 22:31, 1 March 2021 (UTC)[reply]

Yes, it might be too complicated to go in the lead. The natural state variables for the enthalpy of a closed system are entropy, pressure, and chemical constitution. That fits with the idea that enthalpy is naturally measured with controlled (conveniently constant, i.e. isobaric) pressure. For Gibbs free energy, the natural state variables are temperature, pressure, and chemical constitution. As you say, constant-volume calorimetry is suitable for measuring internal energy change. A simple and intuitive way to measure enthalpy is, as you say, isobaric calorimetry, with strict thermal insulation; it controls pressure but lets temperature rise or fall because of the reaction; I have an idea that this is why enthalpy of reaction is considered more useful? I have an idea that this is part of the reason for the term 'heat of reaction'? I defer to your judgement.Chjoaygame (talk) 01:37, 3 March 2021 (UTC)[reply]

In a nutshell, explicitly, in the following, the words "and temperature" are misleading at least, and, unless qualified, are wrong and should be deleted.

In chemistry, the heat absorbed by a chemical reaction at constant pressure and temperature is of interest and is equal to the enthalpy change for the reaction, known as the enthalpy of reaction.

The initial temperature is fixed, so as to define the standard state, but fixing also the temperature of the final state would make it refer to Gibbs energy, not enthalpy.Chjoaygame (talk) 06:25, 5 March 2021 (UTC)[reply]

My apology. I ought to have stuck with my deference to Editor Dirac66's judgment. I remain puzzled.Chjoaygame (talk) 16:23, 6 March 2021 (UTC)[reply]

Thank you. I have been thinking it over also and I believe that most reported ΔH values refer to a specific and therefore constant temperature. For example, if a thermochemical calculations using Hess's law has no heating term (C_pΔT), then all the values are at the same temperature which must be constant. It is true that heat is usually evolved or absorbed in the calorimetric measurement, but the final value determined and reported after correcting for this heat will refer to a specific temperature. Dirac66 (talk) 16:56, 6 March 2021 (UTC)[reply]

I remain puzzled. I feel confident you will be right, because I know your work here. But I still remain puzzled. What exactly is a "correction", and why? Chjoaygame (talk) 19:25, 6 March 2021 (UTC)[reply]

Trying to sort it out, subject to correction:

I think enthalpy of reaction is not a simple enthalpy change in a simple thermodynamic process, as I mistakenly assumed. I think it is a difference between standard enthalpies of two systems of interest. The relation between these two notions is not too simple.Chjoaygame (talk) 19:40, 6 March 2021 (UTC)[reply]

Beginning to understand. We are talking about a compound thermodynamic system, not a simple thermodynamic system.

An initial simple view is as follows.

Initially, we have two simple thermodynamic dynamic systems isolated from one another, and, for the sake of simplicity, from the surroundings. They have each been separately prepared in the standard state of common standard temperature and pressure.

Then one of them is to be regarded as the system of interest and the other as its surroundings, while jointly they are isolated from the rest of the world. The partition between them is removed. The initial system now has a new volume, the sum of the separate volumes. The chemical reaction takes place. The 'intrinsic' energy of the new system is the sum of the 'intrinsic' energies of the separate initial systems. There has been no transfer between the systems and the rest of the world. It is the private business of the new system to choose its temperature and pressure.

Then the new system suffers interaction with the rest of the world. It is no longer allowed to choose its own temperature and pressure. It is forced to gain or lose energy as work by adjusting its volume to reach the standard pressure, and to adjust its temperature by gaining or losing energy as heat to reach the standard temperature. The work is forgotten, but the heat transfer is measured and reported.

Is that right?

A more practical view might be as follows.

The task could be done by use of three containers, and some surroundings with various devices, called "the rest of the world". The outer container has three compartments. One, 'outer', compartment is the "joint surroundings", of measured and controlled and unchanging temperature and pressure, controlled by devices in "the rest of the world", that add or remove energy as heat and work. Inside the outer container is a joint 'reaction' container with variable walls, including one that separates two the initial unreacted systems in their separate 'inner' compartments. They have each chosen their separate equilibrium volumes, and are initially at the standard temperature and pressure. The separating wall is removed. The measured reported quantity is of that of the heat needed from the rest of the world that surrounds the "joint surroundings" needed to keep the "joint surroundings" at the standard temperature and pressure. It seems to need to be assumed that this will not reverse the reaction or take it half way back.

I think there are two mechanisms at work here. One is the expansion of the initial 'system' of interest to fill the final volume and of the initial 'surroundings' also to fill the final volume. This would relate to enthalpy of mixing. The other is transfer of binding energies to other forms in the chemical reaction. A lot for me to get my head around.

Thoughts?Chjoaygame (talk) 23:14, 6 March 2021 (UTC)[reply]

See new section below.

Enthalpy of reaction versus heat of reaction

In order to clarify why the phrase "at constant temperature" is used to describe reaction enthalpies, I think it would be useful to compare the meanings of the two phrases "standard enthalpy of reaction" and "standard heat of reaction", say for the case of combustion where we all know that the temperature is not really constant. Enthalpy is a state function so the value of ΔH depends only on the initial and final states. The standard enthalpy of reaction is defined as ΔH between standard states of reactants and products at the same temperature, usually 298.15, even if the real process involves a combustion generating temperatures of 2000 K or so followed by transfer of heat to the surroundings with cooling back to 298 K. The value of ΔH is said to be at constant temperature because it is the same AS IF the real reaction occurred at constant temperature.

For heat on the other hand this is not true since heat is a path-dependent function. As an extreme example we can cite adiabatic combustion, in which the system is attains a high temperature (say 2000 K) due to the reaction and no heat is allowed to escape to the surroundings, so the real heat of reaction is zero. However this is not the standard heat of reaction. The standard heat of reaction is still defined as ΔH between standard states of reactants and products at 298 K. For combustion this usually corresponds to the heat of the real combustion with increase of temperature PLUS the subsequent loss of heat to the surroundings. And thermochemical calculations involve standard heats of reaction rather than real heats with increase of temperature. — Preceding unsigned comment added by ‎ Dirac66 (talk • contribs) 03:50, 8 March 2021 (UTC)[reply]

Thank you for that. Would it be ok to put into the article some such words as you have just written above? "The standard heat of reaction is [still] defined as ΔH between standard states of reactants and products at 298.15 K [and atmospheric pressure]." I am inclined to think that it demands too much of the reader to ask him to nut out for himself that "The value of ΔH is said to be at constant temperature because it is the same AS IF the real reaction occurred at constant temperature.".

In a way, heat of reaction is a 'latent' heat, with a part analogy with latent heat of phase change: the elementary chemical constitution does not change. Also there might be a heat of mixing if the reactants could be mixed without reacting. Also, it seems to be implied that the return to standard conditions would not reverse the reaction or take it half way back.

Why do they report ΔH, rather than ΔG ? Or do they report both? Or sometimes just ΔG instead? Or even ΔU ? What is the difference for chemistry?

Am I asking too many questions? One might say that this article is primarily about enthalpy as such. Would it be appropriate to indicate why enthalpy is chosen as an energy function to describe chemical reactions?Chjoaygame (talk) 09:04, 8 March 2021 (UTC)[reply]

I will try to answer what I can. First paragraph: it would certainly be ok to put a version of my comments in the article. They need some more editing so I haven't put them in yet, but if you want to insert a version, please go ahead. I think it would be best in the Heat of reaction section; the point is too subtle for the lead.

Second paragraph: Sorry, I don't understand your analogies to latent heat and heat of mixing. But I agree the return to standard conditions is assumed not to reverse the reaction. The initial state is reactants, and the final state is products at the same conditions.

Third and fourth paragraphs: They report ΔH, ΔG or ΔU depending on what is being discussed. ΔH is useful for reactions open to the atmosphere as opposed to ΔU for a sealed container. And ΔH is simpler to understand than ΔG since it does not involve the mysterious entropy. Its measurement by calorimetry is (or seems) also simpler to understand. Dirac66 (talk) 01:20, 10 March 2021 (UTC)[reply]

Thank you for your valuable thoughts.

My part analogy with latent heat of phase change is that a chemical reaction and a phase change are primarily processes internal to the joint system. Primarily they are rearrangements of the constituents of the joint system. Not, for example, primarily simply a spontaneous expansion that does thermodynamic work, nor addition of a selected chemical constituent. The energy changes are primarily attributable to redistribution of energy between microscopic kinetic and various potential forms, especially changes of constituent bonding energy. Mixing is like this too when the expansion of the system of interest is accounted for.

I think the assumption of stability deserves explicit mention. Logically, one would like to choose standard conditions at which the final products are stable, and one would impose those conditions on the initially separate reactant systems.

Entropy is mysterious because traditionalists want to keep it so, by enforcing their nineteenth century 'disorder' dogma, and making the 'ergodic' hypothesis seem like a primary physical fact, putting mathematical convenience ahead of physics. Entropy is an extensive variable, a fact that is obscured by the 'disorder' dogma.

The Gibbs energy has perhaps a claim to relevance here because it is automatically specified with the initial and final temperatures equal. Change of entropy is naturally accounted for by heat transfer. Enthalpy emphasises thermodynamic work.

Instead of 'at constant pressure and temperature', I would like to see something such as 'when reactants and products are reduced to standard temperature and pressure conditions', including in the lead. Yes, wordy, but I think clearer.

Behind my concerns here is that the article isn't well structured to show how the natural state variables for enthalpy are entropy, pressure, and chemical constitution. Gradually, I would like see that remedied.Chjoaygame (talk) 21:47, 10 March 2021 (UTC)[reply]

OK, I have now rewritten the paragraph in the lead. I noticed that Atkins and de Paula avoid saying that T and p are actually constant, and just say instead that T and p are both in the standard state (and therefore equal) in the initial and final states. So I have based my rewrite on that source. Dirac66 (talk) 00:34, 21 March 2021 (UTC)[reply]

relevance of temperature

I refer to this statement in the lead of the article: "The temperature does not have to be specified, ..."

One relevant text that is cited in this article writes on page 66 of its 3rd edition (1999):

Also, the enthalpy change in a reaction varies with the temperature at which the process occurs.

...

Enthalpy changes depend somewhat on the temperature at which the process occurs.^[1]

The IUPAC definition of enthalpy rests on a prior definition of internal energy for a reference state. Without specification of temperature, such a reference state is fully defined by its specified chemical constitution, pressure, and volume. The temperature in such a state is uniquely determined by those specified quantities. This leaves no latitude for a further specification of temperature for the state. I may remark that chemical constitution, pressure, and volume are defined without reference to such characteristically thermodynamic quantities as energy, entropy, and temperature.

For fixed chemical constitution (closed system), for the pressure and volume to change so as to determine a state other than the reference state, the change in internal energy is uniquely determined by changes in pressure and volume.

For fixed chemical constitution, the internal energy, however, can be considered in thermodynamics as a function of entropy and volume, its other natural variables, without specification of pressure or temperature. When a body is allowed to expand against a surrounding pressure slowly enough to allow neglect or precise measurement of the kinetic energy of motion that it imparts to its surroundings, and when the walls are impermeable to conduction and radiation of heat, then its loss of internal energy is at fixed entropy and entirely accounted for only by volume change, as the body loses energy by doing thermodynamic work on its surroundings; pressure and temperature are not specified in such an account.

^ Laidler, Keith J.; Meiser, John H. (1999). Physical Chemistry (3 ed.). Boston: Houghton Mifflin. p. 66. ISBN 0-395-91848-0.

Chjoaygame (talk) 23:17, 2 October 2021 (UTC)[reply]

The statement in question was added on 9 September 2021 by editor Klaus Schmidt-Rohr as part of a major edit. I agree with you that this statement is true, but I am dubious about whether is it is actually useful to include it in the introduction of this article.

It is true because of the phase rule which determines the number of variables needed to specify the state of a thermodynamic system. If enough other variables are given then all other variables including the temperature are fixed and don't need to be separately specified. For example pressure, volume and composition are usually enough to determine all other properties including both temperature and enthalpy, as you have said.

However I believe this statement is potentially confusing to some readers because it may (incorrectly) suggest that enthalpy does not actually vary with temperature, which is not true. So I think it should be deleted by the introduction at least. It could be moved further down with the addition of a careful explanation of what it does and does not imply. However I do not really see what purpose it serves to understanding the concept of enthalpy, even if as you say it is a true statement. Dirac66 (talk) 00:29, 3 October 2021 (UTC)[reply]

Thank you for your comment. I have to say, I haven't yet figured out exactly the best way to express the situation. I haven't yet fully thought through all the dependencies and implications. What I wrote above was intended to provide some starting points.

The question of 'whether enthalpy varies with temperature' is dependent on what other variables are specified, and on whether one is talking merely about the entropy change in some actual particular specified thermodynamic process of interest, or about entropy difference between two sets of standard states, which can require a more complicated calculation, and may be determined without actually conducting the reaction of interest.

(A trivial point. I think it safer to distinguish 'the lead' — sometimes spelt 'lede' here, from 'the introduction'. The lead is more a summary. The 'introduction' is often enough a section of the body of the article, headed with the word 'Introduction', more discursive than the lead.)Chjoaygame (talk) 14:41, 3 October 2021 (UTC)[reply]

In other words, the IUPAC definition of enthalpy is in terms that do not explicitly make it a function of its thermodynamically natural variables. How to deal with this?Chjoaygame (talk) 21:47, 3 October 2021 (UTC)[reply]

First, WP:MOSLEAD considers that both 'lead' and 'introduction' are synonyms which refer to the section before the TOC and the first heading. I prefer to call this the 'intro' or 'introduction', although I realize that some articles do also have a section named 'Introduction' even if this is contrary to the MOS.

In any case, I think most editors agree that the lead or introduction should be kept simple. For Enthalpy, I think it should be based on the textbook (and IUPAC) definition H = U + pV, as in the section Definition. This is aimed at readers who don't know anything about enthalpy. Almost all of them will not know the thermodynamic definition of entropy either, so the intro (or lead) should not mention entropy. Other more complex definitions and properties can be dealt with later in the article.

As for temperature dependence, I think the simplest equation is the first one used in most textbooks, namely dH/dT or more exactly (∂H/∂T)_P = C_P. Natural variables can be dealt with much later in the article, with a link to Thermodynamic potential#Natural variables, and there we can explain why entropy and pressure are "natural variables" and what that term means. Dirac66 (talk) 01:50, 5 October 2021 (UTC)[reply]

Yes.Chjoaygame (talk) 07:28, 5 October 2021 (UTC)[reply]

I have now removed the false statement "The temperature does not have to be specified" and replaced it in the lead by a statement that H varies with T, citing your source by Laidler and Meiser - thank you. Also I have added dH = C_P dT to the section Other expressions. I still have to add the point about natural variables. Dirac66 (talk) 01:24, 26 October 2021 (UTC)[reply]

I think it worth saying here that thermodyamics takes into account various ways of specifying thermodynamic state. This is because thermodynamics is largely about its characteristic variables entropy, temperature, and internal energy. Without use of or reference to those characteristically thermodynamic variables, the condition of a body in its own state of internal thermodynamic equilibrium can be fully specified solely in terms of ordinary physical variables such as pressure, volume, and mole number. Such a specification is not of too much interest for thermodynamics, but it is still physically valid and reliable. There is a wide enough variety of ways of introducing characteristically thermodynamic variables into formulas or equations that describe functional relations between relevant variables. For example, the for a closed system, formula

H(S,P)

indicates that the enthalpy is being considered as a function of the entropy and pressure of the body. That formula is a specification that does not refer explicitly to temperature. In that formulation, if the temperature of the body changes, it is expressed without explicitly referring to temperature; the change is exactly and fully reflected in a change in entropy and/or pressure. Changes in such variables are not causative in a unique and direct physical sense. They just reflect changes of the body's state without saying how the change was physically caused, for example by exposure of the body to rubbing.

If one wants to express causality in terms of a particular formulation, one needs to specify the relevant variables explicitly. For example, to say that

H=U+PV

is, under suitably specified conditions, true as a mathematical relation, but neither

H

nor

U

there is explicitly expressed in terms of its natural variables; both quantities are determined by conventions with respect to standard states, which can be fully specified by ordinary physical variables, without use of characteristically thermodynamic variables. What does it mean to say that the internal internal energy of a closed system is a function of state? Is that referring to a mathematical function? If so, a function of what variables? It is customary to choose to say that

U

is naturally a function of entropy and volume

U(S,V)

? If we do not provide such suitable specification, we condemn ourselves to confusion and bafflement. A mathematical function has a domain and a range.Chjoaygame (talk) 05:08, 26 October 2021 (UTC)[reply]

One could say that to determine H of a substance or ΔH of any reaction, T must be specified either explicitly by giving a value of T, or implicitly by giving values of other variables sufficient to determine the state and therefore T. But I think that is unnecessarily complicated; it is much simpler to just say that these quantities depend on T, and then it should be obvious to most readers that a ΔH at 25°C cannot be assumed equal to ΔH for the same reaction at 100°C so T must be considered. In principle as you say it can be considered either explicitly or implicitly, but in practice most of us find it easier to consider T explicitly in most problems. Dirac66 (talk) 01:10, 27 October 2021 (UTC)[reply]

The Gibbs free energy has temperature explicitly as one of its natural variables, leaving entropy implicit. Enthalpy has entropy as one of its natural variables, accounting for temperature implicitly. In both enthalpy and Gibbs free energy, pressure is a natural variable; volume changes are left implicit, because they are not usually measured. Enthalpy is a more or less traditional way of talking about 'heat content', a natural way of referring to 'heat of reaction', very often measured by calorimetry, and, historically, one of the very roots of thermodynamics. For studying the course of processes, it is convenient to refer to natural variables. For studying bond energies, it is convenient to refer to standard states. They specified by temperature and pressure, which may or may not be natural variables.

I am inclined to think that this deserves expression in the article. I think it isn't mere academic pedantry. For readers who are not too familiar with the topic, I think it will help understand why enthalpy is the customarily chosen variable for studying bond energies and suchlike.Chjoaygame (talk) 19:39, 27 October 2021 (UTC)[reply]

This point could be added to the subsection Other expressions#Characteristic functions which already mentions natural variables twice. Perhaps that subsection should be renamed Characteristic functions and natural variables. Dirac66 (talk) 21:28, 27 October 2021 (UTC)[reply]

I don't want to do too much editing of this article, lest my doing so should detract from the unity of style of the article. So I prefer to leave it to your judgement.Chjoaygame (talk) 22:35, 27 October 2021 (UTC)[reply]

etymology

I have made some efforts at finding Clausius' mind, but I haven't quite enough to nail it against all dispute. Though Gillispie offers 'transformation' as an English translation of Clausius, I think Liddell & Scott's 'turning' is a better rendering of the Greek word. They don't mention 'transformation' in their medium-length entry for the word. Wikipedia is not obliged to be uniform between articles. With external quoting, there is a danger of Wikipedia generating circular errors, that may be hard to undo. If you concur, I would like to follow Liddell & Scott on this.Chjoaygame (talk) 00:31, 26 November 2021 (UTC)[reply]

I think that any detailed discussion of the translation of τροπή belongs not here but in the Entropy article, where I notice that the History section (though not the Etymology section) mentions both 'transformation' and 'turning' and cites the appropriate authors. For this article on enthalpy, we could just insert the alternate translation at the end of the sentence. Perhaps we could say Entropy uses the Greek word τροπή (tropē) meaning transformation or turning. The reader who wants to know the authors of both translations can consult the Entropy article. Dirac66 (talk) 18:49, 1 December 2021 (UTC)[reply]

Good thinking.Chjoaygame (talk) 02:38, 2 December 2021 (UTC)[reply]

Done. Dirac66 (talk) 00:50, 12 December 2021 (UTC)[reply]

Real gas enthalpy...

The blurb:

"The enthalpy of an ideal gas is independent of its pressure or volume, and depends only on its temperature, which correlates to its thermal energy. Real gases at common temperatures and pressures often closely approximate this behavior, which simplifies practical thermodynamic design and analysis."

... can't be right... look at the psychrometric chart... the RH very much affects the enthalpy.

https://i.imgur.com/xoEQCAM.png

Enthalpy for air at 66 F, 100% RH is exactly the same as air at 95 F, 20% RH.

In fact, for conditioning air for buildings, the majority of the energy expended to reduce air enthalpy is via lowering the humidity (condensing it out of the air).

76.30.103.137 (talk) 22:29, 16 November 2024 (UTC)[reply]

[1] Laidler, Keith J.; Meiser, John H. (1982). Physical Chemistry. Benjamin/Cummings. p. 53. ISBN 978-0-8053-5682-3.

[2] Laidler, Keith J.; Meiser, John H. (1999). Physical Chemistry (3 ed.). Boston: Houghton Mifflin. p. 66. ISBN 0-395-91848-0.

[1]

[1]