Talk:Standard enthalpy of reaction

Article is too complicated
Im at a high level of understanding of Chemistry in my opinion, im studying and understanding AS Chemistry, and I don't understand it. What level is this article aimed at? Because to be honest, there won't be many degree graduates reading this! Sorry! Medscin 13:38, 28 January 2006 (UTC)

This is the definition of standard enthalpy change that I learned in high school and again in first quarter freshmen chemistry in college. I don't think this is too complicated. For a basic understanding, readers may use the introductory text, and if they want a mathematical explanation, they can use the general example. However, adding a specific, concrete example may make it more clear. Also, this article should probably be merged with enthalpy, as it is very short by itself. User:carhas0


 * Merging or redirecting with enthalpy would be a mistake. Enthalpy is a more general term that may refer to chemical reactions, heat transfer, convective flow, and other things besides  chemical reactions.  I am removing the merge tag. The article right now is a stub and could use expansion and an example to make it better for the general reader, but it shouldn't be merged.  Flying Jazz 03:14, 13 February 2006 (UTC)

I'M NOT SURE IF THIS ARTICLE IS CORRECT! The standard enthalpy of reaction does not always involve 1 mole of reactants, it depends upon the reaction equation. It is the standard enthalpy of FORMATION that involves 1 mole. This is according to my OCR A2 textbook! As quoted in the textbook (this obviously does not involve 1 mole of reactants): 2H2(g)+O2(g) --> 2H2O(l) ΔHθr=-572KJ mol-1

I find it confusing that the index of summation, B, in the ΔHr⊖ equation also appears in the formula for the generic chemical reaction. I suggest either changing the index to a letter that does not explicitly appear in the chemical formula or using integer indices in the chemical formula: −v1 C1 + ... + −vn Cn → −vn+1 Cn+1 + ... + −vm Cm — Preceding unsigned comment added by 97.114.94.219 (talk) 21:45, 6 April 2014 (UTC)

Nuclear reactions
Is this concept applicable to nuclear reactions? If so and even if not so should be specified in article how and why (not).--188.26.22.131 (talk) 15:24, 27 June 2014 (UTC)

Vandalism
Page got seriously vandalized. No idea how to fix it.189.225.50.104 (talk) 00:20, 25 October 2015 (UTC)


 * I have undone the changes. --DugySK (talk) 13:04, 25 October 2015 (UTC)

needs reliable source
This article lacks a reliable source.Chjoaygame (talk) 16:51, 6 March 2021 (UTC)

According to Atkins & de Paula (eighth edition, 2006, page 29)


 * To avoid a lot of awkward circumlocution, we say that in an exothermic process energy is transferred ‘as heat’ to the surroundings and in an endothermic process energy is transferred ‘as heat’ from the surroundings into the system. ... When an endothermic process takes place in an adiabatic container, it results in a lowering of temperature of the system; an exothermic process results in a rise of temperature.

Coming to basics, they write


 * The energy of a system is its capacity to do work.

In thermodynamics, I think the concrete meaning of this is as follows.

The 'intrinsic' energy of a system is measured as change between a reference state and the state of interest. Of the two states, whichever is of greater energy is allowed to do thermodynamic work on the surroundings until the other state is reached, without any other transfers. The amount of thermodynamic work is given an appropriate sign.

This is the basic definition, but it does not give all necessary and relevant detail. There are many thermodynamic 'intrinsic energies'. Examples are internal energy, enthalpy, Helmholtz free energy, Gibbs free energy, and many more, mostly without proper names. It may be complicated to actually extract work in the desired way. In practice, workarounds may be needed.

This basic definition does not have enough detail for a deduction of a definition of 'heat of reaction'.

They also write on page 30


 * In thermodynamics, the total energy of a system is called its internal energy, U. The internal energy is the total kinetic and potential energy of the molecules in the system (see Comment 1.3 for the deﬁnitions of kinetic and potential energy).[1] We denote by ∆U the change in internal energy when a system changes from an initial state i with internal energy U_i to a ﬁnal state f of internal energy U_f:
 * ∆U = U_f − U_i            [2.1]
 * [1] The internal energy does not include the kinetic energy arising from the motion of the system as a whole, such as its kinetic energy as it accompanies the Earth on its orbit round the Sun.

This seems naughty or unduly vague. It seems to weasel macroscopic internal energy into microscopic energy. Not adequate for our present purpose.

They also write on page 30


 * The distinction between work and heat is made in the surroundings. The fact that a falling weight may stimulate thermal motion in the system is irrelevant to the distinction between heat and work: work is identiﬁed as energy transfer making use of the organized motion of atoms in the surroundings, and heat is identiﬁed as energy transfer making use of thermal motion in the surroundings.

I find this problematic. It doesn't distinguish between isochoric work and thermodynamic work as specified by change in system energy when the system is allowed to do work on the surroundings by exerting a macroscopic force, without any other transfer. I think that thermodynamic work is measured as energy transfer that forces organized motion of atoms in the surroundings, making use of energy that was intrinsic to the system. That is not precisely the same as "energy transfer making use of the organized motion of atoms in the surroundings". Yes, it is importantly required as a matter of principle that the measurement must be made in the surroundings. The original notion of thermodynamic work was that a cylinder of steam expanded and drove a piston rod in the surroundings. That was measured as work in the surroundings. But Atkins & de Paula go on to talk about "energy transfer making use of thermal motion [of atoms] in the surroundings." Carnot didn't think along those lines. In thermodynamics, heat is a mode of transfer of energy, not precisely and explicitly a property of atoms in the surroundings. I think Atkins & de Paula are not reliable at this point.

Still need a reliable source.Chjoaygame (talk) 19:12, 6 March 2021 (UTC)

Notation: E or U
Is there a reason for this article's use of $$E$$ to denote internal energy, while the more common usage is $$U~$$?Chjoaygame (talk) 23:45, 21 March 2021 (UTC)


 * Presumably the first editor to mention internal energy used the symbol E and others followed the precedent. I have just checked the 8 relevant textbooks on my shelves: 5 use U and 3 use E. So I think it is important to tell the reader that both symbols represent the same quantity. Since this article uses E, I will add a note at the first use of E to say that it is often denoted by U.
 * If you want to change all the E to U, that would be all right too, as long as you change the note at the first use of U to say that it is often denoted by E.
 * Strictly speaking, U is probably a better symbol since it denotes specifically INTERNAL energy, as opposed to E which is total energy including kinetic and potential energy relative to the environment. However in practice very few problems involve both quantities, so the use of E for U in thermodynamic problems does little harm if any. Dirac66 (talk) 20:40, 28 March 2021 (UTC)


 * Thank you.Chjoaygame (talk) 04:12, 29 March 2021 (UTC)


 * While notation is under review (strike while the iron is hot), I favour Dirac66's "Strictly speaking, U is probably a better symbol since it denotes specifically INTERNAL energy, as opposed to E which is total energy including kinetic and potential energy relative to the environment."Chjoaygame (talk) 21:15, 9 May 2021 (UTC)


 * I also prefer the use of U to E. My observation is that chemists tend to use U, chemical engineers use U or E, and mechanical engineers and physicists tend to use E. There are similar differences of opinion when it comes to the symbol used for the Helmholtz energy (A versus F). KeeYou Flib (talk) 01:32, 22 June 2021 (UTC)


 * I remember learning that in the past, F was used by some authors for the Helmholtz energy, and by other authors for the Gibbs energy. This from a professor who advised that it is better to use A and G which are less ambiguous. Dirac66 (talk) 02:06, 22 June 2021 (UTC)


 * A wise professor! Since I learned the subject at the knees of chemists, chemical engineers and physicists, I learned this the hard way. Anyway, this is why the IUPAC Green book uses A and G for those quantities (for anyone else who is reading this, see page 48 of https://dev.goldbook.iupac.org/files/pdf/green_book_2ed.pdf). Certainly my own preference, as well as $$\Delta U = q + w $$ (as opposed to $$\Delta U = q - w $$). KeeYou Flib (talk) 03:00, 22 June 2021 (UTC)


 * I prefer $$\Delta U = Q - W $$. I will elaborate my reasons if someone indicates he wants me to.Chjoaygame (talk) 06:05, 22 June 2021 (UTC)


 * In practice both conventions are commonly used, so I can accept either convention. As long as we are consistent within each article and specify which convention is used, which is the case now for this aricle. Readers who want both conventions explained and compared can consult First law of thermodynamics. Dirac66 (talk) 20:23, 22 June 2021 (UTC)

Hess's law
I think that this article should emphasise Hess's law.

In general, the law tells us that the eventual heat of a given reaction can be found through several intermediate steps, and through several pathways involving different chemical reactions, with various respective intermediate, initial, and final conditions such as in solution or as a precipitate, temperatures, and pressures: but that if the overall initial and final conditions are reconciled, then the overall result is the same. The result can be found without performing any actual process of reaction between reagents and products. In order to reconcile the various pathway data, standard conditions, and corresponding thermodynamic quantities, are defined for the respective chemical participants, pure and unmixed, in the overall reaction. Enthalpy is chosen as a suitable thermodynamic quantity for practical and historical reasons. Hess's law was discovered using the concept of heat, before there was a full understanding of the laws of thermodynamics, and before the recognition of the precise character of enthalpy, and was demonstrated by calorimetry.

I think it would help non-expert readers if the article made it clear that Hess's law is essential here, and that the result can be found without performing any actual process of reaction between reagents and products. This may be thought complicated, but I think that, without it, the non-expert reader would be a bit mystified. The complication is the price for the dispelling of the mystery.

There is another quantity that is marginally relevant here, once called the free enthalpy, which I think is now called the Gibbs energy? Check this.Chjoaygame (talk) 05:02, 29 March 2021 (UTC)


 * I suggest adding a section on the evaluation of (standard) reaction enthalpies. However I would not start with Hess's law, which can only be used once we have some measured values for reactions related to the reaction of interest. So first we should mention experimental methods, say 1. combustion calorimetry for complete reactions and 2.the van't Hoff equation for reactions with a measurable equilibrium. Then we can present 3.Hess's law as a method to determine ΔH of a reaction without measurements on that reaction directly. Of course each of these methods is discussed in its own article(s), so in this article we can just present them briefly and add wikilinks for the reader who wants more details.


 * As for free enthalpy, the IUPAC Gold Book definition on Gibbs energy mentions that it was formerly called free energy OR free enthalpy. Also in French, the term "enthalpie libre" is still commonly used, and the French Wikipedia article is called "Enthalpie libre". Dirac66 (talk) 14:43, 29 March 2021 (UTC)


 * A unifying idea is that the standard enthalpy of a pure substance is stoichiometrically referred to the standard enthalpies of its elementary constituents.Chjoaygame (talk) 22:55, 29 March 2021 (UTC)


 * I have now added a section briefly presenting the 3 methods above for evaluating ΔH, with wikilinks to more detailed discussions. Also I mentioned the phrase "free enthalpy" in the article on G. Dirac66 (talk) 02:06, 2 April 2021 (UTC)


 * Great.Chjoaygame (talk) 03:23, 2 April 2021 (UTC)


 * The standard enthalpy of an element is set arbitrarily as a reference point.


 * In general, a physical state can in principle be specified/defined without any thermodynamic concept except the minus–one-th law, that there exist bodies in their own states of internal thermodynamic equilibrium (Callen's Postulate I). Only ordinary physical measurements. Mass (in moles if you like), pressure, volume. None of such thermodynamic quantities as temperature, intrinsic energy, intrinsic entropy.


 * But we want the standard state to be specified at a standard temperature and pressure. We bring the body of the element, with measured mass, to our chosen standard temperature measured with a thermometer. We bring the body to the chosen standard pressure. We measure the volume. We now have specified the standard state by non-thermodynamic physical measurements. From an ordinary physical, non-thermodynamic point of view, it is just incidental and superfluous or redundant that we happen to know its temperature, a thermodynamic quantity.


 * Now we start our thermodynamics. Choose our independent state variables: at least one extensive variable (such as mole number, intrinsic energy, intrinsic entropy) and at least one thermodynamic variable (such as temperature, intrinsic energy, intrinsic entropy), plus the remaining necessary non-thermodynamic variables (such as pressure, volume). Choose our dependent thermodynamic extensive function of state (intrinsic energy, intrinsic entropy).


 * We have now chosen versions of our extensive thermodynamic quantities, intrinsic energy and intrinsic entropy. For example
 * $$H = H(S,P,\{N_j\})~.$$


 * The intrinsic energy might be internal energy, enthalpy, Helmholtz function, Gibbs energy, other unnamed possibilities indicated by e.g. Callen.


 * The intrinsic entropy will be a corresponding variable or function of state. It is customary not to explicitly name the chosen intrinsic entropy with a specific name and symbol as is done for the chosen intrinsic energy. But in principle it deserves a specific name, for it has its own corresponding set of state variables.


 * Euler relations such as


 * $$H=U+PV$$


 * have to be used with regard to the relevant chosen thermodynamic variables of state and functions of state and Legendre transform constants of integration that express the above choices. For example, $$H = H(S,P,\{N_j\})$$ and $$U = U(S,V,\{N_j\})$$ are functions of different arguments and cannot be related for general values by simple identities. Meanwhile, $$P$$, $$V$$, $$PV$$ and $$\{N_j\}$$ are ordinary physical quantities and do not have the complications of Legendre transforms.


 * All this is done as setting up a standard reference formalism for the elements.


 * Then we want values for compounds. For this, we need to actually conduct some chemical reactions, preferably ones that go to completion. We then need to separate (unmix by e.g. distillation, precipitation, condensation, evaporation, sublimation, etc.) the products and get them to pure states at standard temperature and pressure.Chjoaygame (talk) 03:23, 2 April 2021 (UTC)

notation
May I make a little suggestion from the peanut gallery?

The usual notation for mole number uses the Greek letter $$\nu$$, typographically like but not the same as the Roman letter $$v$$. The Greek letter $$\nu$$ corresponds to the Roman letter $$n$$, the initial letter of the English word 'number'. I would favour the usual notation. The Roman letter $$v$$ is often used for 'volume' and for 'velocity'.Chjoaygame (talk) 10:13, 20 April 2021 (UTC)

Checking the 10th edition of the cited text Petrucci & Herring, I see that they use the Roman $$v$$ where the more usual symbol is the Greek $$\nu$$. I can only say again that the idea is 'number' as distinct from 'volume' and from 'velocity', and that the Greek $$\nu$$ is more usual and traditional. I would find it regrettable to depart from reasonable and natural custom and tradition because of an idiosyncratic source.Chjoaygame (talk) 11:08, 20 April 2021 (UTC)

Further checking. I have now looked at the 11th edition of the cited text. It has moved to the use of the Greek letter $$\nu$$. I think this favours our doing the same.Chjoaygame (talk) 12:07, 20 April 2021 (UTC)


 * Good suggestion. Done. Dirac66 (talk) 19:45, 20 April 2021 (UTC)

What do you think about


 * $$\nu_{\text {A}} \text {A} + \nu_{\,\text {B}} \text {B} ~+ ~... \rightarrow \nu_{\,\text {P}} \text {P} + \nu_{\text {Q}} \text {Q} ~+ ~...$$

?Chjoaygame (talk) 21:48, 20 April 2021 (UTC)


 * Thank you. I actually did try to write the chemical equation in math format earlier today, but I couldn't do it properly so abandoned the effort without saving. Now I have copied yours into the article, and added math format for the various ΔH. I think we now have everything in math format which should be. Dirac66 (talk) 01:35, 21 April 2021 (UTC)


 * The way you had it was elegant and neat, and perhaps better. It is questionable whether the LaTeX $$\text {math}$$ formatting is easier or preferable. The spacing of $$\nu_{\text {B}}$$ or $$\nu_{~\text {B}}$$ or $$\nu_{\,\text {B}}$$ is questionable.


 * Just very lately there has been an operating system "improvement" aka 'update' that has damaged the rendering of the $$\mathrm {math}$$ version on my Android tablet. A hyperenthusiastic editor who shall be nameless lately has been busy with a thing that may be closely related. This is beyond me. I guess there may be more coming along soon?Chjoaygame (talk) 06:41, 21 April 2021 (UTC)


 * The "improvement" damages more than the math format.Chjoaygame (talk) 09:09, 21 April 2021 (UTC)

language
I am not familiar with customs in the world of enthalpy of reaction, and so what follows is just my ordinary language intuition and feeling.

When an ordinary language speaker says "when ..." or "where ...", he usually means 'on the occasion of, or at the place of, the actual happening of ...'. When a mathematician says the same, he is likely to mean 'meaning in the present written exposition that ...'.

When an ordinary language speaker says "change", he is usually referring to a particular single individual process. When an ordinary language speaker says "at constant temperature and pressure", he usually means "at maintained unchanging temperature and pressure". In the present account of enthalpy of reaction, "change at constant temperature and pressure" means what an ordinary language speaker would understand as 'increment or difference reduced to standard states'. Many relevant reactions do not actually happen at maintained unchanging temperature and pressure.

Enthalpy of reaction is a constructed quantity derived from consideration of a sequence of ingeniously selected measurements of actual processes. I think it hardly ever refers to an actual single individual process, though it might ostensibly seem to do so.

I feel it would be a generous concession to the ordinary language Wikipedia reader to write, instead of 'change at constant temperature and pressure', such expressions as 'difference reduced to standard states'. I feel it would help the reader to keep in mind how enthalpy of reaction is derived or constructed. I don't know how others would feel about that.Chjoaygame (talk) 10:21, 22 April 2021 (UTC)

The article says "The heat of a reaction depends upon the conditions under which the reactions are carried out." I think that standard enthalpy of reaction doesn't depend on the conditions under which the reactions actually occur. Instead, think it depends on the standard states to which the reagents and products are reduced. I think it would help the ordinary language Wikipedia reader to say so.Chjoaygame (talk) 10:38, 22 April 2021 (UTC)


 * First reaction: I think the phrase "enthalpy change ΔH at constant T and p" can suggest two things to the ordinary reader, one true and one false! The true suggestion is that IF the reaction does occur at strictly unchanging T and p, the heat due to the reaction will in fact be equal to ΔH. This is often true for example for cellular respiration, because T and p in a small biological cell are maintained constant by a whole organism and/or a much larger environment. The false suggestion is that laboratory measurements of ΔH MUST be made at constant T and p. This is often false, for example in bomb calorimetry where the system experiences very high T and p before the heat is expelled to the environment. However because H is a state function, the value of ΔH is the same as if T and p were really constant, provided of course that the system does return to the initial T and p at the end of the experiment.
 * I will think about how to reword the article more clearly. Dirac66 (talk) 19:51, 22 April 2021 (UTC)


 * Thank you for your reply.


 * This is not easy. As above, my first working draft suggestion for new wording is something such as 'enthalpy difference reduced to standard states'.


 * Subtle questions arise. The standard states require that each reagent and each product species be separately in its standard state, a stringent requirement. That might be hard to fit with a metabolic process in a cell, that might not even be alive at the standard state. I don't know how mammalian cells go at 25°C. And, considering the requirement for separation, how would the heats of solution and suchlike affect the numbers?Chjoaygame (talk) 08:37, 23 April 2021 (UTC)


 * But standard states are not necessarily at 25°C; they can be defined at any T and so must be specified. 25°C is the standard temperature at which the values of reaction enthalpies and many other properties are often listed, but the list should always indicate the temperature in question. Dirac66 (talk) 20:19, 23 April 2021 (UTC)


 * I think that we have covered the ground and that you will come up with a good wording.Chjoaygame (talk) 21:54, 23 April 2021 (UTC)


 * I see your new edit. Initially, I would like to propose a small change to it, removing the words "in a system". Changing "The standard enthalpy of reaction (denoted $$\Delta H_{r}^\ominus$$ or $$\Delta H_{reaction}^\ominus$$) is the enthalpy change that occurs in a system when matter is transformed by a given chemical reaction, from reactants in their standard states to products in their standard states" to 'The standard enthalpy of reaction (denoted $$\Delta H_{r}^\ominus$$ or $$\Delta H_{reaction}^\ominus$$) is the enthalpy change that occurs when matter is transformed by a given chemical reaction, from reactants in their standard states to products in their standard states.'


 * My reason is that the standardization procedure involves many processes, and many systems, and that removal of the words 'in a system' does not detract from the intelligibility of the sentence; indeed, I think it makes it easier to read and clearer. Reactants must start as respective pure substances before they are brought together into a single system in which a process of chemical reaction can occur. After reaction, they must be separated into respective pure substances. All the pure-substance systems must be brought to their respective standard states. One might propose that all this occurs in a single system, suitably defined; I think that would not make it easier to understand.


 * If that is agreed, I would like to go further, to remove the word 'change' and replace it with one of the words 'difference' or 'increment'. Some more adjustments to the wording would likely follow from that.


 * My concern is that the standard enthalpy of reaction is a highly engineered quantity, and that the reader deserves fair warning of that. I think it perhaps confusing to suggest that it may be conceived as manifest in a simple process.


 * Here is a possibility that I would prefer. 'The standard enthalpy of reaction (denoted $$\Delta H_{\text {r}}^\ominus$$ or $$\Delta H_{\text {reaction}}^\ominus$$) is the difference between total reactant and total product enthalpies due to a chemical reaction, with reactants and products brought to their standard states.' What do you think of that?Chjoaygame (talk) 21:50, 24 April 2021 (UTC)


 * I can agree as far as the comma, but I would prefer to maintain my phrase at the end in order to specify that the standard state is necessary only for the initial reactant state and the final product state. "Brought to" is not very clear. So how about your phrase to the comma and mine afterwards, giving 'The standard enthalpy of reaction (denoted $$\Delta H_{\text {r}}^\ominus$$ or $$\Delta H_{\text {reaction}}^\ominus$$) is the difference between total reactant and total product enthalpies due to a chemical reaction, from reactants in their standard states to products in their standard states.' Dirac66 (talk) 01:37, 25 April 2021 (UTC)


 * Great. Yes, your part after the comma is better.Chjoaygame (talk) 08:55, 25 April 2021 (UTC)

Coefficients and numbers
In this article, the letter "nu" is used to represent the stoichiometric coefficients but mistakenly refers to these as the stoichiometric numbers (which are positive for products and negative for reactants). Usually the "nu" symbol is reserved for stoichometric numbers. See here: https://goldbook.iupac.org/terms/view/S06025 and also https://chem.libretexts.org/Bookshelves/Inorganic_Chemistry/Modules_and_Websites_(Inorganic_Chemistry)/Chemical_Reactions/Stoichiometry_and_Balancing_Reactions#:~:text=The%20stoichiometric%20coefficient%20is%20the,product%20sides%20of%20the%20equation. Anyone mind if I change the "nus" to n_i and fix the text? I can also use the coefficients and write the enthalpy of reaction as a single sum. KeeYou Flib (talk) 04:18, 5 May 2021 (UTC)


 * I think the article is more readable for non-experts with two separate sums for products and reactants, which means using positive-only coefficients. I agree that the present equation should be described as having stoichiometric coefficients rather than stoichiometric numbers. As for the notation (nu vs. n), I think different authors use different symbols and not everyone follows IUPAC. We could change all the nu's to n's if we are careful to change every occurrence of nu in both equations and text. Dirac66 (talk) 00:25, 6 May 2021 (UTC)


 * That makes sense! Will do. Also will leave the nus alone for now. KeeYou Flib (talk) 14:42, 6 May 2021 (UTC)


 * I agree.Chjoaygame (talk) 15:20, 8 May 2021 (UTC)

? neglecting ...
The article lead says "neglecting the heat of mixing of reagents and products or assuming ideal solutions involved".

I don't understand why that has to be said? It seems to me that the respective species are pure. They are presumably solid, liquid, or gas, not in solution, according to the chosen standard temperature and pressure. The pure respective reactants are mixed (and perhaps dissolved), then the reaction occurs, then the pure respective products are separated. For this, with the relevant calculations, there is no neglecting? Perhaps I misunderstand?Chjoaygame (talk) 16:41, 8 May 2021 (UTC)
 * Don't forget that standard states are also chosen for ions in solution, choosing the enthalpy of formation of H+(aq) at a concentration of exactly 1.00M to be zero. KeeYou Flib (talk) 21:59, 8 May 2021 (UTC)


 * Subtle. I suppose that it is not possible to prepare a solution of pure H+(aq) at a concentration exactly 1.00M. So that's why one would want to neglect the heat of mixing and to assume ideal solutions. So article should articulate, as you have done, that some reactants cannot be prepared in pure form, but require some carrier or receptacle, such as a solvent, and perhaps a counter-charged ion, and even an adjustment for ionic strength, or for pH .Chjoaygame (talk) 06:50, 9 May 2021 (UTC)


 * Yes. Actually the assumption of an ideal solution is a red herring here and I mean to fix that part. The standard states for solutions use actual, real solutions and so if the solute concentrations were exactly 1 M (I erred when I wrote 1.00 M above) before and after reaction, no heat of mixing is neglected. It's the deviations from the standard states resulting from concentration variations that force one to either treat solutions as ideal or nonideal, which is done through the activity coefficients.


 * There are also other choices of standard states. For example, chemical engineers prefer to use 1 molal concentrations for certain applications. Biochemists use a standard state where the H+ concentration is exactly $$1 \times 10^{-7}$$ M. The standard state article actually needs some work because it doesn't address the differences between these conventions (it looks to me like it was written by a chemical engineer), and eventually I'll pitch in there as well. KeeYou Flib (talk) 20:07, 9 May 2021 (UTC)


 * If the initial and final standard states belong to "pure" substances, I don't understand how energy of mixing can be neglected without error? (Assuming a lot about the carrier or receptacle of a "pure" substance.) Various reactants and products may demand respectively different carriers or receptacles? How to articulate such things in the article without making it too complicated?Chjoaygame (talk) 21:27, 9 May 2021 (UTC)


 * It is indeed complicated, and difficult to explain simply in a Wikipedia article or here in this forum for that matter. I've been teaching this material to engineers and chemists for 30 years and I'm coming up short in this explanation, but I'll try. The initial and final states are indeed pure substances whenever possible, but for solutes in solution we are not talking about pure substances anymore and the standard states are defined for solutes in solution (an even for individual ions, using the convention mentioned earlier). Those values are established using calorimetric and/or electrochemical methods at fixed concentration and need not assume anything in particular about whether a solution is ideal. Now, it is true that in some fields - here I am particularly thinking of chemical engineering - they prefer to back-calculate the experimental enthalpies of formation against a hypothetical "infinitely dilute" standard state - not completely sensible IMO when the concentrations of the standard states are so high, but it's done so that activity coefficients can be interpreted (and estimated) as being solely due to non-ideal behavior using Debye-Hueckel theory or such, which is defensible. Here's a classical chemE example: https://www.engineeringtoolbox.com/standard-state-enthalpy-formation-definition-value-Gibbs-free-energy-entropy-molar-heat-capacity-d_1978.html . Biochemists use very dilute concentrations for their standard state and so their choice of a standard concentration of $$10^{-7}$$ M is extremely useful and important in that field. Ultimately the way that the standard state is chosen for solutes with respect to this is completely arbitrary.


 * If all reaction species are solutes and are actually in their standard states of concentration before and after reaction, then the free energy of mixing is exactly zero (because it is calculated relative to the standard values). The entropy of mixing is proportional to the derivative of the free energy of mixing wrt temperature, and so this also vanishes. Since Delta G = Delta H - T Delta S, the enthalpy of mixing vanishes as well. Of course, experimentally the final concentrations are not the same as the initial ones unless the reaction is occurring within an electrochemical cell and voltaic measurement is being carried out. Strictly speaking, this error includes the neglect of mixing but also (and more importantly) neglects the changes in intermolecular forces which happen when the concentrations change. Now, when pure substances are stated as products and reactants and pure substance enthalpies are used to do the calculation, mixing has indeed not yet been taken into account and the enthalpy of mixing has been neglected in the calculation. But if a reaction only involves pure solids, liquids and ideal gases then the enthalpy of mixing is exactly zero, because the pure solids and liquids are unmixed (solids are pure phases to a very high degree of approximation) and ideal gases always form an ideal gas solution. This only changes at very high pressures (say, 30 bar and up). If a reaction involving gases happens at extremely high pressure, then fugacities are used instead of partial pressures in $$K_{eq}$$ and the enthalpy of mixing is no longer zero, although it's almost always quite small compared to the enthalpy contributions from the reaction itself. This is all covered in any reasonably good physical chemistry textbook; see Chang, Chang/Thoman, Atkins / de Paula, Levine, or other similar texts. KeeYou Flib (talk) 14:03, 10 May 2021 (UTC)


 * Thank you for your very careful and helpful post.


 * In order to avoid the difficulties of a fully general statement that covers all contingencies, perhaps the article can make a clean cut, as follows:


 * Primarily, standard enthalpy of reaction is defined for chemical reactions for which all reactants and products can be prepared in their pure forms in standard states; this is relatively straightforward, and is the main topic of this article. Beyond that, however, important reactions involve chemical species that cannot be prepared in their pure forms; for such reactions, standard states need special definitions adapted to the respective reactions; in this article, these cases are covered under specific headings.


 * It seems to me that for reactions involving ions, and other species that cannot be prepared in pure forms, special methods are required, not simply calorimetric. For example, I think some enthalpies of reaction demand measurements of equilibrium constants, which are measured by methods far different from calorimetric? The article could explicitly articulate, and perhaps even explicitly partly exclude, this?Chjoaygame (talk) 20:54, 10 May 2021 (UTC)


 * No, no special methods are required. If for example you measure the enthalpy of hydrolysis of a strong acid calorimetrically, you can use that together with the enthalpy of formation of the aqueous acid to determine the enthalpy of formation of the conjugate base since H+ is defined to be the standard state. In terms of equilibrium constants, they often are not needed. In many cases free energies of reaction are obtained using voltaic cell setups, which are studied as a function of temperature to obtain enthalpies and entropies of reaction. If I understand what you're suggesting (a list of methods for measurement and their limitations), I feel that would result in an overly complicated article, focused on the interests of specialists. I think that the article should be kept as simple as possible, while also being technically correct, so as to be encyclopaedically useful. Does anyone else want to chime in here? KeeYou Flib (talk) 19:36, 11 May 2021 (UTC)


 * Hydrolysis of a strong acid? What reaction is that? Perhaps another word was meant instead of hydrolysis.--86.124.121.86 (talk) 14:49, 18 May 2021 (UTC)


 * The hydrolysis of a strong acid is its reaction with water to form the hydroninum ion and its conjugate base. It's a synonym for acid ionization. See here: https://openstax.org/books/chemistry-2e/pages/14-4-hydrolysis-of-salts?query=hydrolysis&target=%7B%22index%22%3A0%2C%22type%22%3A%22search%22%7D#fs-idm240546096 KeeYou Flib (talk) 15:07, 18 May 2021 (UTC)


 * As Wikipedia always says, I decided to be bold and attempt to address these suggestions and generally improve the article. I hope this will serve. KeeYou Flib (talk) 16:27, 12 May 2021 (UTC)

Now the article mentions mixing corrections only somewhere below the lede, a place not very visible. It should be mentioned very early in article.--178.138.35.88 (talk) 21:26, 17 May 2021 (UTC)


 * This is because the enthalpy of mixing is literally zero for reactions involving ideal gases and /or pure solids and liquids, and is generally (not always!) a small effect for solutions. Featuring the effect as a very important consideration is not quite correct. In addition, I do agree that a more detailed discussion of this belongs in the enthalpy of mixing article (which needs some work!) but I don't see why it belongs in an article primarily focused on explaining the standard enthalpy of reaction. KeeYou Flib (talk) 13:42, 18 May 2021 (UTC)


 * Your first sentence re the zero value of enthalpy of mixing in certain situations should be mentioned in intro. Not mentioning what you say is misleading for those who read only the intro and may get a wrong impression.--86.124.121.86 (talk) 14:09, 18 May 2021 (UTC)


 * That makes sense. KeeYou Flib (talk) 15:07, 18 May 2021 (UTC)

It is worth mentioning that 1M solution for hydrogen ions is a hypothetical dilute ideal solution obtained by extrapolation from infinite dilution, implied by the asymmetric convention for activities.--178.138.35.88 (talk) 21:53, 17 May 2021 (UTC)


 * That is often true, but actually isn't always true. Many authors and practitioners certainly do this, but others use the real solution at 1 M as the standard state and the activity coefficients simply describe deviations from the real solution's properties (which serve as calibration parameters). This is covered in quite a bit of detail in the "Physical Chemistry" textbooks by Atkins and de Paula, Chang and Thoman. This standard state choice is much more difficult to handle theoretically, but much easier to handle experimentally. In any case, I think that a more detailed discussion of this belongs in the standard enthalpy of formation article and not here. KeeYou Flib (talk) 13:42, 18 May 2021 (UTC)


 * Who are these others you say that they use the real 1M solution as the standard state? Are they reliable sources? Are they mentioned by Atkins de Paula, Chang Thoman? This procedure you mention is dubious, both theoretically and experimentally.--86.124.121.86 (talk) 14:36, 18 May 2021 (UTC)


 * Yes, of course they are. And you're entitled to your opinion of course, but we should keep opinions off wikipedia page and focus on things as they are. Dubious? Not at all. The activity coefficients are simply not transferable from convention to convention, and they can't be estimated with Debye-Huckel theory if not related to an ideally dilute solution, but that doesn't mean they are ill-posed or not useful. KeeYou Flib (talk) 15:07, 18 May 2021 (UTC)

notation
Looking at the article on the Van 't Hoff equation, I see the notation

$Δ_{r}H^{⊖}$. It matches Petrucci et al., 11th edition, though not the 10th edition.

This seems like a good idea to me, instead of

$$\Delta H_{\text {r}}^\ominus$$ or $$\Delta H_{\text {reaction}}^\ominus$$.

As a related thing, I think it preferable to use the $$$$ LaTeX format because it exactly matches the equations, and is, in a sense, more systematic, as against the

format, that uses a different font and is distracting if not confusing for the reader.

So I think it would be preferable to write

$$\Delta_{\text {r}}H^\ominus$$ or $$\Delta_{\text {reaction}}H^\ominus$$.

Your thoughts? Chjoaygame (talk) 06:26, 9 May 2021 (UTC)
 * I think these symbols were put in the article some time ago using older tags, and I do agree that they should be updated to LaTeX format.


 * Just so everyone knows - the use of "r" after the delta is a notation specifically designed to denote the change in molar enthalpy due to reaction. It's used in all modern physical chemistry textbooks and some general chemistry texts as well, including Petrucci 11 (a great book btw). It's a notation specific to chemists, though. Mechanical and chemical engineers also work with enthalpies of reaction and they tend to put the "r" after H, not after the Delta. Anyway, I made that change on the Van t'Hoff page and there's no harm in doing this here as well, but definitely use $$\Delta_{\text {r}}H^\ominus$$ and not $$\Delta_{\text {reaction}}H^\ominus$$. KeeYou Flib (talk)


 * In making these edits I noticed that the subscript "r" is already used to denote one specific reactant, so I chose to go with $$\Delta_{\text {rxn}}H^\ominus$$ for the enthalpy of reaction so that "r" didn't mean two different things. I think it works. KeeYou Flib (talk) 16:22, 12 May 2021 (UTC)


 * I see that some changes were made recently to the "products" and reactants" notation I introduced. The problem is that they are not consistent throughout the article. Perhaps an editor was interrupted before finishing? KeeYou Flib (talk) 02:53, 18 June 2021 (UTC)


 * So, I went back and reverted some of those changes so that the notation is the same as that used at the beginning of the article. If anyone has a better idea notationwise, let's work it out here. KeeYou Flib (talk) 18:57, 18 June 2021 (UTC)


 * I think it is mostly OK now. The main reason for my changes of June 15 was that the previous version (June 3) used lower-case p for both pressure and products, which was confusing. However I see that this has been corrected in your latest version which consistently uses P for pressure and p for product. So I will accept your latest version - except for one sentence.
 * The one sentence which I believe is wrong says that "For simplicity we assume below that the stoichiometric coefficients are unity", in the section Heat of reaction at constant volume and at constant pressure. Actually the quantities in most of this section are not identified as molar quantities; for example Eproducts just means the internal energy of the total internal energy of all the products, no matter how many moles there are. Molar standard heats of formation are finally introduced in the final equation of the section, where they are in fact multiplied by stoichiometric coefficients which are not necessarily unity. So I will delete this sentence only, and accept the rest of your edit. Dirac66 (talk) 20:57, 18 June 2021 (UTC)


 * Sounds good - yes, that makes sense. KeeYou Flib (talk) 02:32, 19 June 2021 (UTC)