Talk:Carbonic acid/Archives/2020

Pure Carbonic Acid
From the article:  It is not possible to obtain pure carbonic acid. Why? -- The Anome 09:28, 29 Mar 2004 (UTC)


 * The previous sentence indicates what I believe is the basis for this statement: "exists in an equilibrium with water and carbon dioxide". If you were to somehow obtain a quantity of pure carbonic acid, some of it would immediately and spontaneously break into water and carbon dioxide and it wouldn't be pure any more. Bryan 09:34, 29 Mar 2004 (UTC)


 * If it were pure, then it wouldn't be carbonic acid; it would be hydrogen carbonate. Vacuum c 23:07, Feb 12, 2005 (UTC)

This web page indicates differently so I updated the article http://www.newton.dep.anl.gov/askasci/chem99/chem99661.htm Rjstott 09:37, 29 Mar 2004 (UTC)


 * Pure carbonic acid, O=C(OH)2 DOES exist and has been isolated. When I was in college, many decades ago, we were taught that it did NOT exist as an isolatable species, but existed in solution. (You could, no doubt, find (out of date) references expressing the same incorrect information.) Recently, I became aware (via the journal Science, probably a Oct'14 issue) that it has been known to exist for some time. There will be confusion on this point since a great many of us were taught it doesn't exist. This probably is worthy of inclusion here. Also, the solution properties compose the overwhelming preponderance of information simply because isolation as the pure compound is much more difficult (and much more recent).
 * I also wonder if the expression [H2CO3]/[CO2] should be explained? Obviously those who have been taught (and remember) what square brackets mean in chemistry should have no trouble with this, but I wonder if an encylopedia article shouldn't be more user friendly? There are several issues about it that are problematical: 1. Equilibrium constants are APPROXIMATED by molar concentrations, the actual equilibrium should be expressed in terms of the Thermodyanamic activity rather than the concentration. 2. The expression may be confusing since the definition of the equilibrium constant is activity of products divided by activity of reactants (with suitable exponents) and water is clearly one of the reactants yet is missing from the expression. (As it should be since its activity is 1.0 (when it is the only solvent for dilute solutions).) Its absence might be worthy of comment. & 3. No mention is made about the dissociation into the ions as part of the equilibrium. (Again while this is a first order simplification, roughly 20% of the acid is dissociated in pure water and hence is significant enought to be worthy of note). The equilibrium in pure water comprises Both the hydration and the two possible dissociation reactions. This system comes to equilibrium with a thousandth of the CO2 present as the acid, 20% of that is dissociated to the HOC(O)O(-) anion and a thousandth of that is dissociated to CO3(=), according to the equilibrium constants provided. I don't recall if the pure compound exhibits any self-dissociation in the solid state (it is a solid under the conditions mentioned in the Science article, IIR).173.189.72.93 (talk) 23:27, 18 November 2014 (UTC)
 * Could we please have a proper reference to the Science article: authors, title, volume, issue, and page numbers? Dirac66 (talk) 03:12, 20 November 2014 (UTC)
 * OK, I think I have found it now with Google: Science 31 October 2014 Vol. 346 no. 6209 pp. 544-545, Clarifying the structure of carbonic acid, by Götz Bucher and Wolfram Sander. Please confirm if this is the correct article. Dirac66 (talk) 03:53, 20 November 2014 (UTC)
 * Yes, https://doi.org/10.1126/science.1260117 contians information on the structure. Lpd-Lbr (talk) 15:25, 16 September 2019 (UTC)

Carbonic Acid / Water / Air Equilibrium Notes
Since one of the major components of earth's atmosphere is carbon dioxide, the equilibrium CO2(g) <-> CO2(aq) in combination with CO2(aq) + H2O <-> H2CO3 is important in many solutions that are in contact with air. It is worth noticing that the value of the equilibrium constant of the first dissociation reaction that is usually given (pK=6.36), is really the combined constant of the equilibria CO2(aq) + H2O <-> H2CO3 and H2CO3 <-> H+ + HCO3-. Under normal conditions only a small fraction of the dissolved carbon dioxide reacts with water to form carbonic acid (source: http://www.thuisexperimenteren.nl/science/carbonaatkinetiek/Carbondioxide%20in%20water%20equilibrium.doc, http://www.chem.usu.edu/~sbialkow/Classes/3650/Carbonate/Carbonic%20Acid.html). Therefore the true value of the first equilibrium constant of carbonic acid is approximately 2.6E-4 (pK=3.58). However, as long as the distinction between dissolved CO2 and carbonic acid is irrelevant, the use of the greater pK-value is no problem, but the overall reaction concerned is CO2(aq) + H2O <-> H+ + HCO3-.

Another important quantity involved in the calculation of dilute aqueous solutions in equilibrium with air is the Henry coefficient of CO2, which describes the first equilibrium CO2(g) <-> CO2(aq). At 293 K it is about 25.6 kg atm/mol (Edwards, Newman & Prausnitz, AIChE J., 1975, 21(2), 248-). This value corresponds very well with the observations that about 39 mmol of CO2 dissolves in 1 L of pure water of 293 K (under pure CO2 at 1 atm., final pH 3.9) and that the pH of water in equilibrium with air (0.03% CO2 by volume) is 5.7.

--Watje22 13:46, 3 August 2005 (UTC)

Carbonic Acid Contradiction
The table says it exists only in solution, but the article links to a page which says that it DOES exist as an indipendent substance. Isn't that unconsistent?--Army1987 21:03, 15 August 2005 (UTC)
 * Both are wrong! The press release claims that carbonic acid has been isolated in the gas phase, but goes on to say that the preparation of gaseous H2CO3 is an objective for researchers into interstellar space... I am sceptical about gaseous H2CO3, the article cited describes only computer calculations, not actual preparation. However, carbonic acid can be prepared as solid, crystalline etherate (like a hydrate but with dimethyl ether), melting point -47 &deg;C . An etherate is not H2CO3, but it is certainly not a solution either. Physchim62 01:46, 16 August 2005 (UTC)

Carbonic Acid Density
The density listed for carbonic acid is shown as 1.0 g/cm3 in a dilute solution. However, the density listed for water is shown as 1,000 kg/m3, which is the same thing in different units. Shouldn't these entries use the same units? Also, I believe I read somewhere that carbonic acid is slightly less dense than water, so it tends to be a bit more common near the surface, but Googling only seems to turn up copies of the Wikipedia information. Can anyone verify density of carbonic acid? Thanks. --70.20.161.132 23:59, 10 October 2006 (UTC)
 * This is indeed the same figure in two different units: g/cm3 is the unit which is more commonly used. There is no significant difference in the density of carbonic acid and water, at least at normal pressures of carbon dioxide. Physchim62 (talk) 08:13, 12 October 2006 (UTC)

Acidity
The Ka value for this acid is diffrent in my chemistry text book. Ka1=4.4*10^-7 ka2=4.7*10^-11
 * Ka2 is (roughly) correct, Ka1 is false because of the equilibrium between carbonic acid and dissolved carbon dioxide, as discussed in the article. Physchim62 (talk) 17:39, 13 November 2006 (UTC)

Rate law for H2CO3 dissociation
What rate law is assumed for the rate constants given for carbonic acid dissociation / CO2 hydration? I assume the rate law is first order in both directions - i.e. rate = k[CO2] or rate = k[H2CO3]. However, this should be stated explicitly. Asteen 15:48, 3 October 2007 (UTC)

Rate and equilibrium constants
Could you please add sources for the forward and reverse reaction rate constant values and also the coresponding equilibrium constant value you cite? Thanks, Griffgruff 12:23, 25 October 2007 (UTC)

Disolved CO2 in rain, molecular ratio
I'm trying to understand the removal of co2 from the atmosphere by rain (say over the oceans). Assuming a rain drop is in equilibrium with the surrounding air (380ppm co2), pH 5.7 say, just before it hits the ocean surface, what is the ratio co2 molecules to h2o molecules? If you can provide the weight ratio as well perhaps I can avoid making an error in the calculation. My guess for the molecular ratio is about 1 in 3000. Also in the formation of rain drops is it fair to assume a gradual build up of volume as the drop is form, thus integrating co2 into the raindrop from the beginning of its formation, making the assumption of equilibrium uncontestable. Any thoughts? blackcloak (talk) 04:16, 20 January 2008 (UTC)

Proportion dissolved by atmosphere/liquid by time?
Hi, is there a resource showing the proportion dissolved by altitude (atmosphere, liquid) over time? 69.228.195.34 (talk) 17:48, 31 August 2008 (UTC)

It's not just the insertion that is hard
Keeping one in has become a struggle as well. Too many vested interest, trying to stop sources from being added. They even want to judge what is considered a source. It is starting to look like Atlas shrugged. FX (talk) 11:54, 31 January 2010 (UTC)
 * Clearly a sinister conspiracy against what is good and pure.--Smokefoot (talk) 14:08, 31 January 2010 (UTC)

Strength of carbonic acid
Article states that carbonic acid is a stronger acid than acetic acid and formic acid, even though the latter two have higher Ka values... Error? —Preceding unsigned comment added by 174.88.19.86 (talk) 00:19, 28 January 2011 (UTC)

"Remark" with approximate formula for [H+]
This concerns the "remark" leading to H+ concentration at equilibrium. The author gives a result using certainly a too crude approximation which gives very different values for "carbonate concentrations" in aqueous solutions. In fact, a more rigorous relationship, leading to H+ concentration, must also be used in this case. After some substitutions and taking account of the different equilibrium constants, a term in H+ concentration remains in a secondary degree equation (in H+ ). When solving this secondary degree equation, one is facing an indetermination (of the form zero over zero). The classical Hopital rule may thus be used in order to solve this indetermination. As a result, it is found that when carbonate concentration is ten to the minus four the corresponding pH is seven and when carbonate concentration is ten to the minus six, the corresponding pH is six. So, the carbonate concentrations are found to be of several orders of magnitudes greater than the values proposed by the author. Beranay — Preceding unsigned comment added by Beranay (talk • contribs) 16:27, 15 February 2011 (UTC)
 * I just clarified/corrected this. Hope you like it :) OneAhead (talk) 17:44, 17 August 2018 (UTC)

Reference
Here is a Reference: http://antoine.frostburg.edu/chem/senese/101/solutions/faq/dissolving-gases.shtml Tideflat (talk) 05:20, 1 February 2012 (UTC)

Values in "Acidity of Carbonic Acid" are wrong
After having my edits "undone" I apparently need to go further than editing some bad math. First of all the Ka1 value is incorrect. The pKa1 value IS correct but if you do the math you will find the value of Ka1 is 5.0x10^-7 (3 orders of magnitude different than the one listed!!!!!). This changes ALL of the values for [HCO3-] AND pH in the data table listed. If certain editors could do math then we could get correct information out there! — Preceding unsigned comment added by 128.118.16.157 (talk) 21:37, 28 February 2012 (UTC)

I thought organic compounds contain of H and Cs
In article it says that "Carbonic acid is the inorganic compound" but it have hydrogens and carbon, is that a mistake or I don't know something? — Preceding unsigned comment added by 71.105.186.28 (talk) 23:50, 7 April 2012 (UTC)
 * Usually organic compounds are considered to have C-H bonds, but there are many arguments about the definition. Not worth going there.--Smokefoot (talk) 01:59, 8 April 2012 (UTC)

Information about boiling and melting point of pure carbonic acid
Why there is no information about melting and boiling point of this compound? It is very interesting question: Is pure carbonic acid solid at room temperature? It is written taht this compund can be isolated in gas phase... — Preceding unsigned comment added by 79.191.60.250 (talk) 21:20, 10 December 2012 (UTC)

Wrong hydration equilibrium constant?
The current hydration equilibrium constant of 1.7e-3 has no reference. A published paper ("CO2 system hydration and dehydration kinetics and the equilibrium CO2/H2CO3 ratio in aqueous NaCl solution", Soli and Byrne, Marine Chemistry, vol.78, p.65-73, 2002) states a significantly different value: 1/848 = 1.18e-3. Has anyone got a reference for the current value or should we just change it? [Post by Tomasvandenooijevaer as per history page]
 * I believe it's quoted as 1/600 (=1.67e-3) in D.M. Kern, The hydration of carbon dioxide, Journal of Chemical Education, 37(1): 14–23, 1960. Having said that, I don't think there is a consensus out there, one way or the other. The value of the eqb const depends on a number of factors, like pressure and salinity, so being precise to 3SFs is probably not possible due to experimental error, and certainly not without stating a whole list of other conditions. Also, the eqb const is just the ratio of the two rate constants, which have each been quoted variously by factors of at least 2 in different references (all apparently at 25C). In my experience, if you can get these figures correct to within the right order of magnitude then you're doing well! That's not to say we shouldn't change the webpage... PS. Please sign your posts :-) Mmitchell10 (talk) 20:39, 11 March 2013 (UTC)
 * Yes, I think salinity (ionic strength) is the key factor. The article by Soli and Byrne says their value is in 0.65 molal NaCl(aq) which is close to the ionic strength of seawater. The value 1/600 is presumably for pure water with no NaCl. I think both values should be quoted in the article as the pure-water value is a fundamental property of CO2(aq), but the value in NaCl(aq) is of great environmental importance.
 * Another reference for the value in pure water is Housecroft and Sharpe, Inorganic Chemistry, 2nd ed, Prentice-Pearson-Hall 2005, p.368 which gives K ≈ 1.7e-3. Also note that the ratio of the two first acidity constants in the Acidity section is Ka(app)/Ka1 = [H2CO3]/[H2CO3*] ≈ [H2CO3]/[CO2] = Kh = 4.6e-7/2.5e-4 = 1.8e-3, using notation and values from the acidity section. Dirac66 (talk) 21:33, 11 March 2013 (UTC)
 * Thanks, that clears things up a bit - should have spoken to you years ago! I've made the change Mmitchell10 (talk) 20:27, 14 March 2013 (UTC)

Flash Point
The infobox lists the flash point as 169.787 °C, this is an error. Flash point is defined by the temperature of a liquid where the vapor above it will form a flammable mixture with air. Carbonic acid and the CO2/H2O which it is in equilibrium with are not flammable at all ratios in air. Therefore, carbonic acid has no flash point or respective flash point hazard. 169.232.128.66 (talk) —Preceding undated comment added 23:01, 12 April 2013 (UTC)


 * I have moved this new comment to the end of the talk page. It appears that you are correct, as this material data safety sheet for CO2 lists the flash point as "Not applicable" (in section 5). So I will go ahead and remove this mistaken information from the infobox. Dirac66 (talk) 23:46, 12 April 2013 (UTC)

Recently
To the Wikipedia Editor: Never use the words "recent" nor "recently" as they quickly become incorrect with time and require someone to edit them. This article had "has reportedly been recently disproved" but today it was five years old. Instead use "has, in 2011, reportedly been disproved" which will be correct forever. Nick Beeson (talk) 12:03, 16 February 2016 (UTC)

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Carbonic acid pure
Practical: if we think carbonic acid is as energy. This mass materials vacuum in very high-pressure block. In center thermodynamics system. From center cool to hightemprature valves and pipes join to gather it's generate infinite energy and lights. Prashant Nanda 15:06, 19 October 2016 (UTC) — Preceding unsigned comment added by Twisindia (talk • contribs)

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hi
hi নাফিস আহমেদ নিলয় (talk) 02:34, 28 February 2017 (UTC)

untitled
\:The acid dissociation constant is defined as a ratio of concentrations, two in the numerator of the fraction and one in the denominator. It usually (but not necessarily) has the units of mol/L (or mol/dm3, which is the same but longer to type). Physchim62 (talk) 15:16, 2 September 2007 (UTC)

Acidity
The contents of this section are badly garbled.

There are 4 reactions to consider.
 * (1) First dissociation of carbonic acid in solution: H2CO3 HCO3- + H+. pKa1 applies.
 * (2) Second dissociation in solution: HCO3- + H+ CO32- + H+. pKa2 applies.
 * (3) Dissolution of CO2 in water. CO2 (g) CO2 (aq) Henry's law applies
 * (4) Hydration of dissolved CO2. CO2 (aq) → H2CO3 (aq) The concentration of water is effectively constant when this reaction occurs in dilute solutions.

Reaction (1) is effectively complete at pH > 7. Reaction (4) is slow in the absence of an enzyme, but may be assumed to go to completion when the system is at equilibrium. Therefore only (2) and (3) need to be considered for solutions at biological pH if it is assumed that equilibrium conditions are achieved. Petergans (talk) 16:09, 20 October 2020 (UTC)

General revision
The basis for this revision was the separation of biological and non-biological material, resulting in much greater clarity of presentation in both areas. Some irrelevant material has been removed. Petergans (talk) 11:43, 6 November 2020 (UTC)

The pKa values are being edited to and from. The first equilibrium step is quite complex, depending on whether the hydration equilibrium is incorporated or not. 3.6 and 6.3 are reported for the two cases. However the second step is much weaker, pKa2 around 10.3. Both the text and the data box should reflect this. There is a plot that is using the 6.3 and 10.3 for species distribution. — Preceding unsigned comment added by 2A02:AB88:1543:8100:7545:74E1:49BE:6C57 (talk) 01:59, 4 December 2020 (UTC)