Talk:Equilibrium constant

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[edit] Article replaced

The original article has effectively been merged with chemical equilibrium. The present text is completely new.

See talk:chemical equilibrium for more details.

Petergans 17:43, 16 May 2007 (UTC)

[edit] Untitled

Kc and Kp are also dimensionless, as they are defined properly using activities of the reactants and products which are dimensionless too.

For simple calculations for Kp, dividing by the standard pressure of 1 bar for each component in the ratio brought to any power will always yield a dimensionless result.

Daniel0ng 13:43, 21 April 2007 (UTC) Daniel.

Isn't it K=[(C^n)(D^p)]/[(A^k)(B^m)] ?

I do not know who wrote this but I agree about that. I'm going to correct. Ange Blanc

[edit] units

can any one write something about the units of the reaction constant and how they are generated?

If you learned anything about equilibrium, you should know that equilibrium constants don't have units. No offense meant, but thats the one thing my teacher seemed to think was important about equilibrium, so I guess I have it deeply ingrained. Kr5t 02:41, 29 March 2006 (UTC)

I've been taught otherwise. Equilibrium constants can have units; in fact, they only don't if the equilibrium reaction in question involves the same number of moles on both sides (e.g. A+B->C+D has a Kc with no units, but A+B->4C involves a Kc with units). The units of Kc can be worked out simply by the relevant equation: for instance, in the reaction expressed as A+B->4C, Kc = [C]^4 / [A][B], and therefore its units are mol^2•dm^-6. - (Anonymous), 17 May 2006

Im sorry but anonymous above is actually incorrect. The equilibrium constant is always dimensionless because you devide each term by the standard state, effectively "cancelling the unit." I would refer you to "Atkin's Physical Chemistry", by Peter Atkins (Professor of Physical Chemistry at the University of Oxford) which i believe is the standard text for 1st year chemistry undergradutes. - (Markus), 28 July 2006

I agree with Anonymous, it is exactly what I was taught too. To get the units correctly you must write the units as if writing the normal reduced-over-oxidized equation... If none are squared (or cubed, whatever) then K is dimensionless. If one of the reduced is squared, the overall unit is mol/L3. If one of the oxidised is squared, the unit is (mol/L3)-1. —Preceding unsigned comment added by 195.158.102.80 (talk) 16:32, 26 April 2008 (UTC)

An equilibrium constant can be expressed as being dimensionless (Ko, or simply K): however, it is more usually referred to in terms of Kc or Kp, both of which can have dimensions [1]. This approach avoids having to use activity coefficients or fugacities, which are essential for the dimensionless approach (read Atkins more carefully!) Physchim62 (talk) 13:40, 28 July 2006 (UTC)


-- I agree. Each concentration should be divided by standard concentration, in the same units as the concentration. if the concentrations are in mols/liter (say), and the standard conc is 1 mol/liter, then the division has no numerical consequences, but it DOES get rid of the units. However,if the concentrations happen to be in molecules per cubic Angstrom, if you're a molecular simulation person :), then the standard concentration of 1 mol/liter needs to be included explicitly and its numerical value is something like 1 molecule/1660 cubic angstroms. So this WILL change the numerical value of K (compared to omitting it), and it will, again, get rid of the units. I think this issue is worth getting right in so widely used a widely-used reference site as Wikipedia! (Unsigned)

Let's be absolutely clear about this. The expression

\ \Delta G^{std} = -RT \ln K

requires that K be a (dimensionless) number because only numbers have logarithms. K can be written as the product of a concentration quotient and an activity coefficient quotient. The two quotients have reciprocal dimensions. If the latter quotient is assumed to have a value of unity the numerical value of the concentration quotient is the same as the equilibrium constant. The device of dividing each concentration by a unit concentration works because the corresponding activity coefficient can be multiplied by the same unit concentration, leaving the activity coefficient quotient with a value of one and a K value which is is unchanged.Petergans 22:24, 24 April 2007 (UTC)

This is, of course, completely true: however, don't forget that the numerical value of K (and of ΔG°) depends on the form in which you write the reaction equation. Physchim62 (talk) 16:41, 26 April 2008 (UTC)
Perhaps it should be added that in biochemistry it is not unknown for a "binding constant" to be written as though it had units. For example, the binding constant for a host-guest complex, HG, can be written as K=\frac{[HG]}{[H][G]} so that apparently K has the dimension M-1 when concentrations are molar. Perhaps this is done to indicate the scale of measurements. Is this the source of the confusion above? Petergans (talk) 20:27, 26 April 2008 (UTC)

[edit] Merge with Chemical equilibrium

This really oughtta merge with Chemical equilibrium, and with solubility equilibrium, and with an added section about solubility product constant, one of t he more special equilibrium constants since it doesn't involve the reactants.—The preceding unsigned comment was added by Daniel0ng (talk • contribs).

Let's discuss about a merge over at Talk:Chemical equilibrium. --HappyCamper 16:46, 24 April 2007 (UTC)

[edit] Graph

The graph in this article is useful however it is slightly misleading. As a student we mostly study Gibb's free energy (and its relation to the Equilibrium constant) with spontaneous reactions of which this graph represents none. It would be nice to show the lower side of this curve too in order to allow for a better perception of the concept.

[edit] Poor applications section

A perfectly sensible tidy up of this section I made was reverted. I've re-inserted the cleaned up wording of this section but left the dubious examples:

For example, in the Haber process for the formation of ammonia, the value of K is around 30 at pressures and temperatures standard for the process. In an equilibrium between two conformers with energy difference 0, the equilibrium constant is 1 and both conformers are present in a 1:1 ratio. When the energy difference increases to 1 kcal/mol, the equilibrium constant at 25 °C becomes around 5 and the concentration of the more stable conformer gets 85%.

Not only is this poorly worded but they are random examples which add nothing i.e. detract from article. Left for discussion.

Similarly the Derivation section is muddled and:

When the equilibrium constant is close to unity and the reaction rates very fast for instance in conformational analysis of molecules, other methods are required for the determination of rate constants for instance by complete lineshape analysis in NMR spectroscopy.

is just bizarre.

What was the scientific reason for reverting this deletion?

84.92.241.186 18:10, 30 April 2007 (UTC)


  • Please consider getting a user account. You have bluntly deleted two sections from this article and I did not have any seconds thoughts reverting your edits as a whole. It is considered vandalism. Also get some manners: my edits in general are not considered dubious, muddled or bizarre (you must be English). I am always prepared to discuss controversial sections. Simple add a {{Fact}} tag and its up to me to provide an explanation and a reference. The first statement that you find dubious is from March or from Vogel so no problem there. The second statement simply explains that the equilibrium can establish itself slow or fast and that special techniques are required when they are fast. V8rik 19:53, 30 April 2007 (UTC)
    • I'm trying not (but failing?) to get sucked into the potentially bottomless pit that is Wikipedia - hence the lack of account. But I take the point that a drive-by edit by an anon user will arouse suspicions (despite WP:AGF). I'm happy to discuss my take on things off-site (nospamATphodgkin.plus.com) but in short: there's nothing wrong with the 1st statement but it is equivalent to "the Haber process is run under conditions that make the Haber process favourable" (gosh!). As for the 2nd statement, I'm afraid the explanation leaves me even more confused and I teach all this stuff! Any way, I've found that the parallel Chemical equilibrium article is more to my tastes, so I'm a bit calmer about things (although still not convinced that Wikipedia will work for this stuff). 84.92.241.186 21:27, 30 April 2007 (UTC) (Random chemistry professor)


  • Thanks for your reply. The image giving free energy vs equilibrium constant is my adaptation of a table in carey/Sundberg Advanced Organic Chemistry Part A Sec. Ed. p 144. The first statement gives an idea how to interpret the image. I have thus far neglected to include the ref but I will solve that. I like the Carey / Sundberg table because it tries to relate different quantities. Also Wiki chemistry should not be about explaining chemistry to chemists but to the general public and these sort of details will specifically help these readers. I understand a brand new chemical equilibrium article is about to float so much is happening but also notice the merge discussion on its talk page. Yes I do too believe Wiki will eventually fail but for other reasons V8rik 20:56, 1 May 2007 (UTC)

[edit] Effect of temperature

This addition is unsatisfactory for various reasons.

1. To quote from chemical equilibrium

"The effect of changing temperature on an equilibrium constant is given by the van’t Hoff equation

\frac {d\ln K} {dT} = \frac{{\Delta H_m}^{\Theta}} {RT^2}

Thus, for exothermic reactions, (ΔH is negative) K decreases with temperature, but for endothermic reactions (ΔH is positive) K increases with temperature. An alternative formulation is

\frac {d\ln K} {d(1/T)} = -\frac{{\Delta H_m}^{\Theta}} {R}"

This is much clearer and more informative. A corollary is that K for an athermic reaction (ΔHo=0, ΔSo positive) is independent of temperature. In this context it should also be stated that ΔHo is itself a function of T (depends on heat capacity), so the the second equation gives an only approximately linear relationsip of log K with 1/T.

2. " The related Nernst equation in electrochemistry gives the difference in electrode potential as a function of redox concentrations.

When molecules on each side of the equilibrium are able to further react in secondary reactions the final product ratio is determined according to the Curtin-Hammett principle."

These statements are irrelvant to temperature change and should be removed.

3. The graph is all wrong as it puports to show free energy as a function of K, but it should be K as a function of T as that's what this section is about. Also the axis labelling is faulty and the caption "More stable isomer" is incomprehensible without the original context. A better graph would be of log K as a function of 1/T (straight line) for various values of ΔHo, such as -20, -10, 0, 10, 20 kJ mol-1 for example.

Petergans 09:04, 15 June 2007 (UTC)

Hi Petergans, I just reintroduced the whole section because like I stated before, I was opposed to its deletion. The title is new and as it appears now creates more confusion. I think that both the Nernst equation and the Curtin-Hammett principle are integral part of equilibrium and should be incorporated either in chemical equilibrium or equilibrium constant. Suggestions please where exactly it should go. The plot is there simply to give an idea what sort of energies are involved in equililibria between two isomers. I will redo the plot and I am open to suggestions (reversing the axes is okay with me.) I disagree with the faulty axis labeling, exactly what is wrong? V8rik 20:17, 15 June 2007 (UTC)

  1. I agree with the idea that there should be a section on the effect of temperature changes, but the current content is vague and lacking basic material.
  2. The Nernst equation is not currently included in chemical equilibrium but is referenced in the See also list. I suggest you put a question on that discussion page to the effect : should there be a section on redox equilibria? My own view is that it's rather complicated, particularly in relation to redox electrodes, and best left as a reference. It certainly does not belong where it is at present.
  3. The Curtin-Hammett principle is about kinetics and as such has no place in a discussion about equilibria. The first sentence in that article is wrong in this sense: the equilibrium product ratio does not depend on transition state energies, but on the free energy difference between the two products.
  4. When there are two independent variables on an axis it is customary to show them on opposite sides of the plot, like top and bottom for the x-axis. Labelling on the y-axis is confusing as it seems like one quantity is subtracted from another. In any case kJ mol-1 would be sufficient on its own. Note that the notation kJ/mol is deprecated by IUPAC since the plot is of numbers, that is, of ΔG/kJ mol-1 vs K; number=quantity/unit.

Petergans 08:31, 16 June 2007 (UTC)

  • The reference to Nernst equation should never have been deleted in the first place. A reference is sufficient, no need to duplicate information
  • Your Chinese wall between kinetics and thermodynamics is artificial and without merit. Check out Free-energy relationships to see how interrelated they are.
  • I am not convinced about you criticism regarding the plot. V8rik 20:50, 16 June 2007 (UTC)

You are mistaken. It is not a Chinese wall. Let me put it this way: does the activation energy for a reaction have any bearing on equilibria? Of course it does not. It follows from this that kinetics has no bearing on the composition of a mixture at equilibrium. Petergans 09:58, 17 June 2007 (UTC)

Free-energy relationships are empirical and often inadequate. Actually, kinetics is a deeper level of looking at things than equilibrium thermodynamics. The latter is much more easier and used more often, but it can be derived if there is a complete knowledge of the former, in the same way that thermodynamics can be derived from Statistical mechanics. The chemical potentials of all equilibrium (and non-equilibrium) species can be calculated from a complete knowledge of the potential energy surfaces, so it is not entirely correct to say activation energy does not have any bearing on equilibrium composition (the very concept of activation energy is itself dubious, as it ignores entropy considerations which can often dominate reactions, particularly biochemical ones). The inter-relation between kinetics (or in a broader sense non-equilibrium thermodynamics) and equilibrium thermodynamics is quite complex, and the differences are great enough to require emphasis. I agree that purely kinetic phenomena should be kept out of this article. However it is often superficial to strictly avoid all references to kinetics or mechanisms when discussing equilibrium. Loom91 08:36, 18 June 2007 (UTC)
  • It is important to emphasize that no kinetics are introduced in this article, kinetics are used where needed. I also want to point out that it is important to first explain things in a simple way with focus on educational value and then discuss the finer points. You cannot outright dismiss Free-energy relationships (please point me to literature). This whole discussion reminds me of an earlier discussion about where a bunch of chemical engineers where about to scrap the entire Rate equation article with the argument that systems in real life are not closed. That really does not help wikipedia V8rik 19:40, 18 June 2007 (UTC)

I think this argument can be very easily resolved. Chemical equilibrium has a lot of stuff about Gibbs energy, so Free-energy relationships can be referenced in the See also section of that article. That is where the reference belongs, not in the body of the article, as they are somewhat peripheral and add nothing of substance to the topic. Petergans 08:05, 19 June 2007 (UTC)


  • I am not going to settle with the see also section. I am against it for the reason that a lot of links are presented without any context at all to the article. I will not insist in having Free-energy relationships included at all and I think that with both oscillating reactions and Curtin-Hammett statement in just two lines in a section of chemical equilibrium already overwhelmed with links, their presence is very modest. V8rik 19:18, 19 June 2007 (UTC)
I support including oscillating reactions, but Curtin-Hammett principle has little relation to equilibrium. I think it's way off the topic. Loom91 20:58, 19 June 2007 (UTC)

[edit] Nomenclature

Please follow standard nomenclature. See e.g. [2].   —DIV (128.250.80.15 (talk) 08:54, 16 March 2008 (UTC))