Talk:pH/Archive 2

parts Hydrogen (pH)

I added "parts hydrogen" to clearly define that pH is a direct measurement of the concentration (parts) of hydrogen ions in a liquid (compared to distilled water). --Lperez2029 13:09, 15 December 2006 (UTC)

It isn't a direct measurment of hydrogen ion concentration. It is definitely a logarithmic scale, not linear. The only other problem being that if by distilled water you really mean pure water, then your explanation will yield a 'divide by zero' error (or efectively a pH of infinity). Pure water contains no hydrogen ions (or any ions at all as no dissociation takes place). By strict definition, pure water cannot have a pH value. It is not even posible to measure the pH of pure water, because, as I noted above, the measurement process itself introduces all of the measurable ions, completely invalidating any result.

The article already contains several theories as to what pH stands for, none of which remotely resemble 'parts hydrogen', and it wouldn't anyway as it is a logarithmic scale, so something involving 'powers' is clearly closer. I B Wright 12:46, 16 December 2006 (UTC)

pH is the symbol for the logarithm of the reciprocal of hydrogen 'ion concentration' in gram atoms (parts) per liter. For example, a pH of 5 indicates a concentration of .00001 or (10 to the power of 5) gram atoms of hydrogen atoms in one liter of solution. Therefore, pH measurement may be expressed as parts (ion potential) hydrogen, or pH (potential hydrogen), either way it is acceptable in definition terms, but "parts" hydrogen is more correct (as would the expression of ppm (parts per million). No edit wars - I'll bring some top chemists into the discussion, and we will see what is more correct for pH definition for the benefit of all, once and for all. Best regards --Lperez2029 00:39, 17 December 2006 (UTC)

I did a bit of internet research to come up with that last paragraph. Instead of getting "top chemists" into the discussion, I'd recommend you ask them for existing scientific, peer-reviewed literature that discusses the nomenclature of pH. My survey is deficient in the fact that it does not consult traditional print or academic sources, but if you can find reputable research that says otherwise, I'd have no problem including it. — Edward Z. Yang^(Talk) 03:35, 17 December 2006 (UTC)

pH literally means negative the log of the hydrogen ion 'activity' in solution; it is commonly written -log[H+] meaning -log of the hydrogen ion concentration. In the common solvent water, a bare hydrogen ion is solvated by the water forming the hydronium ion [H(H2O)]+, so pH is also commonly written -log[H3O+]. For example, pH=7.0 means that the positively charged hydrogen ion concentration in solution is 1x10-7 moles H+/liter, since -log(1x10-7) = 7. Normally, pH would therefore be close in meaning to parts hydrogen (as in ppm) with the understanding that hydrogen means positively charged hydrogen ions in solution; whether the concentration is reported as pH, ppm, g H+/liter, wt%, etc., - moles H+/liter, ie concentration, is the basic chemical term.

There are some fine points here. The 'activity' or 'effective concentration' of H+ is what's measured by a pH meter electrochemically. In dilute solution, the 'activity' and actual concentration of H+ are indistinguishable; activity = concentration. As concentration increases, relative 'activity' decreases due to a chemical thermodynamics concept called chemical potential. Activity < concentration, and concentration H+ is undermeasured by the pH meter because the 'activity' is lessened due to this thermodynamic ionic shielding factor. (The actual H+ concentration could be measured by other means such as titration). So, in this context, one might say that the 'p' has to do with the potential H+ concentration (which is activity) which is less than the actual H+ concentration.

Water does inherently dissociate. Pure water, under 'standard' conditions, dissociates from H2O into hydrated H+ and OH- to the extent of 1x10-7 moles/liter, also called pH 7 or neutral pH, the pH where there is no excess of acid (H+) or base (OH-). Regardless of the early 1900's origins of the use of p in pH, consistent with the meaning of 'p' in descriptions of other chemical equilibria, such as pK, the p of pH nowadays is taken to mean -log.

Now, let the statement speak for itself - as for adding 'parts', I'll leave that up to you gentlemen. Best regards --Lperez2029 22:17, 18 December 2006 (UTC)

You are absolutely correct, and this is already discussed in the article. We have to tread a very fine line between giving correct information and giving "usually correct" information (but more commonly used). If you feel you can clarify the article in this respect, please do so. Personally, I get the feeling that the opening paragraph is a little convoluted.

As for "parts Hydrogen", I would refrain for now. A quick Google for "Parts Hydrogen" reveals that the term is used mostly as a kind of more generalized mole. — Edward Z. Yang^(Talk) 21:28, 21 December 2006 (UTC)

Thanks Edward. --Lperez2029 23:17, 21 December 2006 (UTC)

The opening paragraph is a bit convoluted (see Intro below) and needs attention, as someone already asked a very good question, "is there a separate test to measure alkalinity?" Go ahead Ed, refine the article, I'll back you up. --Lperez2029 13:32, 22 December 2006 (UTC)

It is suggested that pH stands for "pondus Hydrogenii" (=~ weighting of hydrogen-ions). --march 15, 2009 —Preceding unsigned comment added by 85.81.91.41 (talk) 14:28, 15 March 2009 (UTC)

Intro

"pH is a measure of the acidity of a solution in term of activity of hydrogen (H+). "

Wait but pH also measures the alkilinity of a substance. Is there a seperate test for this? Simply south 15:34, 21 December 2006 (UTC)

pH is also the electrochemical measurement of the Alkalinity of a solution. Alkalinity is closely related to the acid neutralizing capacity (ANC) of a solution. --Lperez2029 18:00, 21 December 2006 (UTC)

To more specifically answer your question; no, there is no separate test necessary to measure alkalinity in a solution, the test may be performed utilising a pH meter or litmus paper. The introduction on the article should read more along the lines pH is a measure of the acidity or alkalinity of a solution in terms of activity of hydrogen (parts Hydrogen). For example; 4= Acidity - - 7pH=pure water - - 12=Alkalinity (or base) --Lperez2029 13:20, 22 December 2006 (UTC)

As has been noted elsewhere, such an introduction would be incorrect as pure water does not have a pH value of 7, or any pH value, as it contains no ions. I B Wright 11:32, 12 January 2007 (UTC)

Alkalinity is not the opposite of acidity. Alkalinity is a measure of the buffering capacity against the addition of H+. And yes, there is a separate test for this (titration). pH is a measure of the H+ activity. period. From this 'basicity' can be determined, if you know the pKw (ie pOH = pKw - pH)(the pKw shouldn't be assumed to be 14 - it is rare to have standard conditions), but there is no way to determine alkalinity from pH as you need to take into account a wide variety of other ions that contribute to alkalinity (ie At = OH- + HCO3- + CO3-- + B(OH)3 + ...). It is likely that there is a way to measure basicity ( ie [OH-]), though it is presumably much easier to measure acidity ([H+]). Piyrwq 16:43, 12 January 2007 (UTC)

Anybody else have a problem with this statement: "strictly speaking, there is no such thing as the H⁺ ion"? There certainly is such a thing as the H⁺ ion. It is called a proton and exists discretely in the in the state of matter known as plasma if nowhere else. If nobody objects, I would like to change this statement to read: "in aqueous solution the H⁺ ion does not exist independently but is actually the third hydrogen atom in H₃O⁺, called the hydronium ion." Thoughts? Objections? Schlemazl 21:53, 13 February 2007 (UTC)

Defintitely an improvement; maybe a little absolute though, there is probably one free aqueous H⁺ out there somewhere. Maybe something like "in aqueous solutions H⁺ ions bind to water molecules forming hydronium ions, or H₃O⁺." Piyrwq 02:41, 14 February 2007 (UTC)

Good point. I'm pretty sure there is a free H⁺ out there because I lost mine the other day while I was bathing. I got out of the tub, dried myself off, and dismantled the drain to see if it got caught in the P-trap. No luck. If anybody finds a free aqueous H⁺ out there please hold on to it until you can e-mail me a picture of it so I can try to determine wheter it is mine or not. If nobody finds it in another day or so, I'll go ahead and make the change as suggested by Piyrwq. Otherwise I think we should use my wording unless anybody else can provide evidence of a known free H⁺ in aqueous solution.Schlemazl 18:20, 14 February 2007 (UTC)

Also, I'm being driven a little crazy by the addition of "basicity" or "alkalinity" to the definition of pH. Knowing [H₃O⁺] does not mean you know [OH-] (basicity). If you have the literature resources at hand, measured temperature, and know the specific nature of other ions in solution you may be able to calculate basicity from the pH, but one does not define the other. It definitely does not give any insight into alkalinity. For example, it is possible to reduce pH of a solution, through the addition of CO2, while alkalinity remains constant Alkalinity#Addition_of_CO2- alkalinity is absolutely not defined by the log of hydrogen concentration, nor is basicity. These are concepts worthy of much discussion, but are not fitting for the definition of pH. Can we come to some kinda consensus, so that references to basicity and alkalinity can be permanently removed from the definition of pH? Piyrwq 02:41, 14 February 2007 (UTC)

Rewritten

Ok, I have rewritten the whole first paragraph - I hope it is an improvement, but no doubt there is still a lot to be changed. For reference, here is the old version, which I felt was over complicated and too wordy.

pH is a measure of the acidity and the basicity/alkalinity of a solution in terms of activity of hydrogen (H⁺). In aqueous solutions H⁺ ions bind to water molecules forming hydronium ions, or H₃O⁺. For dilute solutions, however, it is convenient to substitute the activity of the hydrogen ions with the concentration, molarity (mol/L) of the hydrogen ions (however, this is not necessarily accurate at higher concentrations ).

In aqueous systems, the hydrogen ion activity is dictated by the dissociation constant of water (K_w = 1.011 × 10⁻¹⁴ M² at 25 °C) and interactions with other ions in solution. Due to this dissociation constant, a neutral solution (hydrogen ion activity equals hydroxide ion activity) has a pH of approximately 7. Aqueous solutions with pH values lower than 7 are considered acidic, while pH values higher than 7 are considered basic.

The concept was introduced by S.P.L. Sørensen in 1909, and is purported to mean pondus hydrogenii in Latin. However, most other sources attribute the name to the French term pouvoir hydrogène. In English, pH can stand for "hydrogen power," "power of hydrogen," or "potential of hydrogen." All of these terms are technically correct.

Conrad.Irwin 10:10, 21 February 2007 (UTC)

This is incorrect. In 2000, Jens Nørby, a Danish biochemist, went back to Sørensen's original papers on the subject: "But Nørby traces the 'p' to a simple mathematical convention of naming variables. (...) But he did not give an explicit reason for choosing 'p' in his original two papers. Sørensen was primarily concerned with determining the H+ concentration electrometrically. His central equation involved values for measurements at two electrodes, which he arbitrarily designated p and q. To develop a standard, he set the non-hydrogen component, Cq, at 1.0 and solved for the hydrogen ion concentration, Cp, or 10–p. The number p he suggested calling p+H."

Source: http://www1.umn.edu/ships/words/pH.htm which is a summary of the original article: Nørby, Jens. 2000. The origin and the meaning of the little p in pH. Trends in the Biochemical Sciences 25:36-37. ($31.50) —Preceding unsigned comment added by 75.74.75.149 (talk) 20:05, 1 February 2009 (UTC)

I like your new intro a lot, but I did make some copyedits to it. I hope you find them satisfactory. A few questions though:

• The original introduction purported that it was impossible for water to truly have a pH of 7. I believe this assertion is correct
• Could you source your mathematical derivation of the symbol pH? It sounds very plausible, but I'd sleep better if it was sourced.

Thanks for stepping up to the plate! — Edward Z. Yang^(Talk) 21:43, 21 February 2007 (UTC)

The new introduction is good on readability, but there are some inaccuracies. First, 7 is not "defined" as neutral, and as Edward Z. Yang says, pure water doesn't have a pH of exactly 7 (maybe it does at some temperature, but not at 25 deg C). Also, although this is an omission of the entire and not just the intro, the pH concept applies as well for solvents other than water, with their own dissociation constants and different "neutral" pH values and different pH ranges. Then, the statement "Because pH is dependent activity, a property which cannot be measured easily or predicted theoretically, it is impossible to give an accurate value for the pH of a solution" is vacuous, because it depends on the degree of accuracy desired. On one extreme, it is also impossible to measure the length of a sheet of paper accurately, but on the other, it is certainly possible to measure pH with a certain degree of accuracy such as +/- 0.01 pH units (note: I'm not an expert on pH measurement so I don't know what is the maximum accuracy attainable in practice). pH is measured every day with no worries, so to imply in the introduction that it is almost impossible to measure is too deprecating IMO. Finally, saying that "the H should more correctly be [H+], standing for concentration of hydrogen ions" is a POV, and seems like a contradiction to the earlier definition that uses activity instead of concentration. --Itub 09:05, 22 February 2007 (UTC)

Playing off Itub's (valid) point of pure water not having a pH of exactly 7, it seems to me that the intro still needs fixed. I will make a point that the pH is not exactly 7 due to the self-ionization of water, and wikilink that to the appropriate page. --ExNoctem 22:02, 2 May 2007 (UTC)

Dimensionless

Regarding "By virtue of its logarithmic nature, pH is a dimensionless quantity." IANACB (I am not a chemist, but..) is it truly because of its logarithmic nature? Wouldn't say, a quantity expressed as "log10 lightyears" still be very much a dimensioned distance? Does the dimensionless of pH perhaps have nothing to do with logarithims and everything to do with the value whose logarithm is being taken? If I'm wrong here, thanks for your patience in considering this. —Preceding unsigned comment added by Dfgriggs (talk • contribs) 18:22, 22 November 2008 (UTC)

You are both right and wrong. Right in the sense that you can only take the logarithm of a pure number, and wrong because the logarithmic quantity, pH, requires that the hydrogen ion activity must be dimensionless also. BTW, you cannot take the logarithm of a quantity like light years, which has the dimension of time, but you can take the logarithm of the number of light years, as the number is equal to the quantity divided by the dimension. Petergans (talk) 18:38, 22 November 2008 (UTC)

What p stands for

The present intro says, "It is unknown what the exact definition of p stands for. Some references suggest the p stands for “Power”[2], others refer to the German word “Potenz” (meaning power in German)[3], still others refer to “potential”."

To be fair, "power", "potenz", & "potential" are all rough synonyms. Ventifax (talk) 00:00, 1 April 2009 (UTC)

____________________________________________________________

pH is the abbreviation of the latin expression pondus Hydrogenii (pondus = pressure, hydrogenium = hydrogen) —Preceding unsigned comment added by 165.228.210.24 (talk) 04:59, 2 April 2009 (UTC)

Contradictory statement within the article

The article states: Most substances have a pH in the range 0 to 14, although extremely acidic or extremely basic substances may have pH less than 0 or greater than 14. An example is acid mine runoff, with a pH = -3.6. Note that this does not translate to a molar concentration of 3981 M.

But it also states that pH is, strictly speaking, a function of the activity of hydrogen ions in solution, not their concentration. And that is acknowledged by references. Now, a careful look at Wikipedia's article on Activity, or preferably to a Physical Chemistry book, will lead you to see that Activity ranges from ZERO to ONE, and is never higher than one, because pure substances have an activity of 1 by definition (it was meant to be analogued to molar fraction, and molar fraction can never be greater than 1 either).

Now pick a calculator and verify the value of -log(1), the result is 0, and that is the lowest pH a substance can ever have. Strictly speaking, pH can never be lower than zero, except when calculated in a less rigurous way, using molarity or molality: but that is incorrect, and should be pointed out. pH lower than zero may be obtained by other ways, maybe (pH-meter?), but it should be clearly specified how was that value obtained, what does it mean (because if it has no meaning, is serves no purpose and shouldn't be used), and emphasis should be put on that strictly speaking, the pH of a solution can never be lower than 0 (because activity ranges from 0 to 1, and therefore -log(act) ranges from 0 to positive infinity). I'll wait for answers regarding this before modifying the article myself. 158.170.51.28 14:50, 27 December 2006 (UTC)

activity expressed in mole fraction ranges from 0 to 1. Activity can be expressed using molarity (as with pH) or molality and can exceed 1. Chemist207.8.130.130 18:27, 28 December 2006 (UTC)

You are confusing the activity coefficient γ, for which 0<γ<1 and α=γ*C where α is the activity and C the concentration in mole/L. You can have α>1 and thus pH<0. Still, I don't understand why pH= -3.6 is not translated to concentration (even effective concentration). How is it measured? and if it doesn't represent the concentration then what does it represent?
Also you can't have a concentration of H+ in an aquatic solution higher than that of water (55.5 M) so you can't even reach pH= -2.0, let alone pH= -3.6. What's the nature of the acid mine runoff? Anyway I don't know of any solution which can reach a density so high that would allow C>100 mole/L. As it is, this example is confusing instead of enlightning.--84.254.22.153 17:42, 14 January 2007 (UTC)

I have looked into the matter and found out that γ can be >1. In fact these values are so high exactly because γ is extremely high and not concentration. I am editing the page accordingly including the reference. Since I am new at this maybe someone can format my edit so that it looks better/it is linked to other parts of wikipedia.--84.254.20.189 17:39, 27 February 2007 (UTC)

Puissance d'Hydrogène

Was added by User:69.140.68.72 and removed by me. A Google search brings up 20 hits. The top four English resources (disregarding Wikipedia's status as first):

• http://limnology.wisc.edu/courses/summer315/handouts/4.Handout.pdf
• http://www.arroyoseco.org/wqparameters.htm
• http://www.swrcb.ca.gov/nps/docs/cwtguidance/3140fs.pdf
• http://iach16.chem.klte.hu/~lenteg/Docs/Oct30.pdf

— Edward Z. Yang^(Talk) 15:03, 19 February 2007 (UTC)

Images

This article needs a nice image to plaster on the top of the document. One can't really photograph pH, though, so we've got to pick something fairly representative. Candidates:

• Image of the pH scale
• Image of a pH indicator

Possible candidates? (add ones you find to the gallery below) — Edward Z. Yang^(Talk) 21:53, 21 February 2007 (UTC)

• 
• 
• 

This page has been vandelised somewhat

Not experienced with this, but i cant seem to delete the vandelism :( —The preceding unsigned comment was added by 129.234.4.76 (talk) 14:50, 24 February 2007 (UTC).

Also there seems to be a discrepancy between tables "Representative pH values[citation needed]" and "pH in body fluids [13]" Apologies if this is in the wrong place but text editing is not my forte. -->Check the Gastric Acid #'s

Semiprotection

I propose we semi-permanently semi-protect this article. The amount of vandalism it receives is absolutely ridiculous. — Edward Z. Yang^(Talk) 22:36, 1 March 2007 (UTC)

Semiprotected. Please yell at me if you disagree. — Edward Z. Yang^(Talk) 17:54, 16 March 2007 (UTC)

Why is this article still protected? Was the vandalism so bad that it has to be protected for 2 months now? Dismas|^(talk) 05:05, 15 May 2007 (UTC)

biological values?

I think it would be interesting to have some biological Ph values for comparison. For example the stomach. Also according to de.wikipedia wine has Ph 4, there is a hole there in the english table. --172.174.140.43 07:47, 4 March 2007 (UTC)

pH-imbalance

I actually wanted to look at the page "pH-imbalance", meaning in the body, but I just got redirected to pH.

Different scales of pH

If somebody feels up to the job, you might want to incorporate the fact that there are three different scales for pH in place, as a recent discussion on the talk page to the article ocean acidification has showed. Hardern 12:01, 13 April 2007 (UTC)

Wow, I had no idea about that. I don't know much about those scales, but from what I have been able to find, they seem to be used only in the context of chemical oceanography. Something worth mentioning in the article, but going into it in a lot of detail might be too much for a general article about pH (we can always have an article on acidity of seawater ;-) ). --Itub 12:12, 13 April 2007 (UTC)

Note: After looking further, I've found that some of these scales are discussed in books that have nothing to do with oceanography. --Itub 12:22, 13 April 2007 (UTC)

There's also a fourth scale, called the NBS or infinite dilution scale. A good source available from google books is this: [1]. --Itub 12:26, 13 April 2007 (UTC)

An excellent reference for this task: [2] Piyrwq 00:38, 16 April 2007 (UTC)

I've now written some text here to try to describe the different pH scales in use in oceanography. My intention is to add this to the pH article in the near future. However, I'm very much not a chemist, and would appreciate comments on it before I do so. Cheers, --Plumbago 15:04, 17 April 2007 (UTC)

I've just added the text that I've been working on over at Talk:Ocean acidification (where interest in different seawater pH scales began). There's also been discussion there about it there, but future discussion should probably move to here now. Hope the added text is OK by everyone here. Cheers, --Plumbago 08:54, 19 April 2007 (UTC)

The text looks very good. I made a small correction, and deleted a line about Ks* that seemed out of place (and debatable?).--Piyrwq 21:16, 20 April 2007 (UTC)

I can't really comment on whether it's a debatable point, not being a chemist and all. The textbook I drew on made a point of mentioning the difficulty of estimating K_S^* as a limitation of the free scale. Certainly, in my own work, I've only ever used the total and seawater scales - I'd never heard of the free scale in oceanography until I read the aforementioned textbook's dissing of it. The free scale sounds like it makes sense in freshwater or weak solutions, but any solutions with significant concentrations of ionic species may have trouble with it. From some of the discussion about seawater pH it transpires that blood has its own pH scale too - perhaps looking at that might illuminate things? I'm sure someone cited a web reference somewhere (perhaps over at Talk:Ocean acidification). Cheers, --Plumbago 08:51, 22 April 2007 (UTC)

pH of lemonade

Two problems with the following:

For example, if one makes a lemonade with a H⁺ concentration of 0.0050 moles per litre, its pH would be:

${\mbox{pH}}_{\mathrm {lemonade} }\approx -\log _{10}{(0.0050)}\approx 2.3$

Firstly, the word lemonade refers to totally different drinks in the UK and USA. The carbonated drink commonly known as lemonade in the UK (and limonade in France) is better known as Sprite in the US, but without the lime flavo(u)r. Secondly, the article does not state whether 0.0050M is a reasonable proton concentration for lemonade - is this a hypothetically concentrated lemonade that nobody would want to drink, or a realistic one? The article on lemons says that pure lemon juice has a pH of 2 to 3, and the article on acetic acid says vinegar has a pH around 2.4. Assuming that lemon juice in lemonade is diluted with about 5 parts water, this would give it a pH 0.8 higher than pure juice - and I would also assume that (American) lemonade is less acidic than vinegar, if sourness is any guide to pH. Mtford 06:36, 10 May 2007 (UTC)

I say, be bold. Change the parameters so they make sense. — Edward Z. Yang^(Talk) 00:41, 11 September 2007 (UTC)

Propose Phidron be merged into this

If it's real, it seems that Phidron would be better as part of this article. A search of Google Scholar did not find any references to phidron. Most of those in Google itself are apparently based on this article. ChemGardener 19:55, 28 May 2007 (UTC)

the small p

"However pH is actually a shorthand for its mathematical approximation: in chemistry a small p is used in place of writing − log10 and the H here represents [H+], the concentration of hydrogen ions."

My understanding was that the use of p as shorthand to mean -log developed from pH, not the other way around, so I'm pulling this unless someone can source it.

Orange32 03:35, 27 August 2007 (UTC)

I am adding one more question to this. Is it usually written as p^H, I mean is 'H' a superscript of p or is it just pH? Elncid (talk) 12:33, 20 December 2007 (UTC)

If it was ever written as p^H, we would write so in the article. --Cubbi (talk) 14:11, 20 December 2007 (UTC)

Effects of pH

This may require a separate article, and I am not an expert (or even proficient) by any means on the following, but the following sections seem to be missing:

• The effect of pH on plants and greenery
• The effect of high and low pH biologically on the human body

Perhaps a separate article may be needed. —Rob (talk) 20:05, 15 October 2007 (UTC)

Hear, hear. —Preceding unsigned comment added by 58.109.98.211 (talk) 06:28, 27 November 2007 (UTC)

So what you're asking for is the effect of acidity and alkalinity on the environment, and living things. Correct? Those articles have the specifics, and then some.64.252.67.170 (talk) 20:55, 18 March 2008 (UTC)

don't understand

I don't understand something from this artical, Im a fifth grader so this question must seem stupid to most of you but the less acidic something is like pure water you would say want to put a fish in it, at least someone told me at a feild trip. So you want to put fish in bleach and household lye? —Preceding unsigned comment added by 76.104.227.88 (talk) 04:23, 29 November 2007 (UTC)

This is a discussion page about the encyclopaedia article, not a tutorial about acidity and biology. Your question is really not relevent. 130.88.117.216 (talk) 10:24, 1 February 2008 (UTC)

Too much talk, not enough information

This article is too confusing for someone who just wants a simple explanation of pH. It needs a better introduction and some explanation in words, not just math. —Preceding unsigned comment added by Amberlianne (talk • contribs) 03:22, 9 March 2008 (UTC)

I wholly agree with this. pH is just a simple acidic calculator, not 10 mathematic equations which have nothing to do with acids and science. Same with Circumference. Koshoes (talk) 18:54, 25 January 2009 (UTC)

Indeed! I have a PhD in Computer Science and informal Chemical education, and I find this article level totally inappropriate. I think the "solution" is to split it into an "Exact definition of pH for Chemical Engineering" and keep "pH" as a high-school grade article which does not make references to "activity" at all. A basic article on pH should follow most biology/chemical 101 text, defining it as -log [H+] using molarity, and giving specific examples of actual pH and its importance to the human body and simple solutions (eg swimming pools, fish tanks, etc)

see for exmaple http://antoine.frostburg.edu/chem/senese/101/acidbase/faq/what-is-pH.shtml

Seeing as everything you've described is already in the article, all I can gather from your suggestion is that you want to put the existing subsection p[H] before the existing subsection pH? I suppose that would make some sense, they would then be in historic order, sort of like we've done with the countless definitions of acids and bases at Acid-base reaction theories --Cubbi (talk) 15:09, 1 May 2009 (UTC)

pH

Is there a simple way for the article writers and editors to state plainly the question, "What is pH?" 70.231.150.141 (talk) 02:11, 19 March 2008 (UTC) In a simple way for any range of high school student down to elementary school student to understand? Thanks. 70.231.150.141 (talk) 02:13, 19 March 2008 (UTC) Wait, I'll post it as best as I can. —Preceding unsigned comment added by 70.231.150.141 (talk) 02:14, 19 March 2008 (UTC)

Discard hydronium?

I think the article would be clearer if we dropped most references to hydronium ions (H₃O⁺) in favour of hydrogen ions (protons, H⁺).

When you consider the solvation, H₃O⁺ isn't all that accurate anyway - it's more like H₉O₄⁺ but this depends on temperature. So H₃O⁺ isn't much more accurate as a description of what protons do in aqueous solution than writing H⁺ (aq). However, writing H⁺ is much simpler and scarcely less accurate.

Ewen (talk) 08:33, 23 March 2008 (UTC)

Sounds good to me. pH is a measure of a formal "hydrogen" and in practice it is defined operationally due to the complexity of the physical chemistry of the solution. --Itub (talk) 11:03, 25 March 2008 (UTC)

How's about this for an intro?

Reworking the introduction section somewhat, how does this sound?

pH is the measure of the acidity or alkalinity of a solution. pH is formally a measure of the activity of dissolved hydrogen ions (H⁺), but for very dilute solutions, the molarity (molar concentration) of H⁺ may be used as a substitute with little loss of accuracy. In solution, hydrogen ions occur as a number of cations including hydronium ions (H₃O⁺).

In pure water at 25°C, the concentration of H⁺ equals the concentration of hydroxide ions (OH^-). This is defined as "neutral" and corresponds to a pH level of 7.0. Solutions in which the concentration of H⁺ exceeds that of OH^- have a pH value lower than 7.0 and are known as acids. Solutions in which OH^- exceeds H⁺ have a pH value greater than 7.0 and are known as bases. Because pH is dependent on ionic activity, a property which cannot be measured easily or fully predicted theoretically, it is difficult to determine an accurate value for the pH of a solution. The pH reading of a solution is usually obtained by comparing unknown solutions to those of known pH, and there are several ways to do so.

The concept of pH was first introduced by Danish chemist S. P. L. Sørensen at the Carlsberg Laboratory in 1909. The name pH has been claimed to have come from any of several sources including: pondus hydrogenii, potentia hydrogenii (Latin), potentiel hydrogène (French), and potential of hydrogen (English).

I've tried to limit the hydronium ion stuff (not least because the hydrogen ion pages points to other hydrates of protons, and implies that "hydrogen ion" is favour by IUPAC), and have broken the text up into several more readable (to me) paragraphs. I've also extended the mention of acids and alkalis, since this seems pretty important to me.

Anyway, I'm no chemist, so may have botched some of the above. Comments welcome. If I don't hear anything at all, I'll replace the current intro with the above. Cheers, --Plumbago (talk) 16:05, 4 April 2008 (UTC)

OK, no comments for now. I'll swap the text above into the main article. --Plumbago (talk) 12:43, 8 April 2008 (UTC)

The assertion that pure water is pH 7 will confuse budding experimentalists who find that their "pure" water is acidic. Perhaps it would be appropriate to mention that pure water exposed to the atmosphere becomes acidic as CO₂ combines with water to form carbonic acid. This is often a point of confusion as folks are experimenting with pH. JonMoulton (talk) 20:36, 8 April 2008 (UTC)

Good point, although perhaps not in the article intro. In this short section of text, my feeling is that this sort of detail would probably just confuse general readers (see some of the comments about simplification in sections above). However, this sort of information would fit very well into the section on measurement though. Cheers, --Plumbago (talk) 08:00, 9 April 2008 (UTC)

Dissociation versus ionization

In text: "But here we define a strong acid as a species which is a much stronger acid than the hydronium (H3O+) ion. In that case the dissociation reaction (strictly HX+H2O↔H3O++X− but simplified as HX↔H++X−) goes to completion, i.e. no unreacted acid remains in solution"

Dissociation ?

Contributions from Brazil

user:pt:lemarlou —Preceding unsigned comment added by 189.68.63.196 (talk) 15:05, 11 April 2008 (UTC)

Gastric acid

There are two tables in this article. One places gastric acid at 1.5-2.0, the second places gastric acid at 0.7. Which is correct? --Arperry (talk) 22:28, 29 July 2008 (UTC)

gastric acid has another two values to add: 0.8 for the acid where it is secreted from the epithelial cells, 1.0-2.0 for the overall acidity of the stomach content.

So my guess is 0.7 would be about the most acid you could get, 1.5-2.0 is the typical range for stomach content, and 1.0-1.5 is probably heartburn.

RandomP (talk) 14:27, 20 August 2008 (UTC)

The tables should be made to agree with each other. If I felt better qualified, I'd make the change myself, but I'm a physics man and do not know these things. --Arperry (talk) 19:18, 10 September 2008 (UTC)

volt

I find it a little sad that both the article and the talk page currently fail to include the string "volt". So, well, at least the latter I can remedy:

At 293.15 K, a pH difference of 1 corresponds to a voltage of .05817 V, or ln(10)×R×293.15 K/F precisely. For other temperatures, modify accordingly.

That much I am able to gather from the definition given on the page. pH meter appears to have slightly more, but could the relationship between pH and voltage on the one hand, and pH and temperature on the other hand, be made clearer to the non-experts among us?

Is it really as simple as saying that a voltmeter will display pH differences, if temperatures are constant, in the ideal case of an infinitesimal test charge being used?

RandomP (talk) 13:51, 18 August 2008 (UTC)

Can pH gradients be used as a lens?

http://en.wikipedia.org/wiki/Talk:PH A recent news article claimed that a pH change of as little as 0.02 units could significantly change the degree of absorption of sound in seawater, especially at low frequencies.[3] Do small pH changes also affect the speed of sound in water? To illustrate, picture a large explosive sitting on the ocean floor, and far above it a net of electrodes establishing a steady pH gradient, or a net of tubes emitting fluid of high/low pH, or a large wrapped isoelectric focusing gel... when the target sub passes overhead, the sound of the blast is refracted by differences in transmission speed, and a significant fraction of the original blast is reconstituted at the distant focal point. Or not? ;) Wnt (talk) 17:55, 3 October 2008 (UTC)

Image Scaling

The pH scale images are too small to interpret from the page. This would render them totally useless in printed versions. Can we get them scaled up, please? --IronMaidenRocks (talk) 19:02, 7 November 2008 (UTC)

pH Chart

Doesn't pH only go down to 0? Why does the chart say -1? Shouldn't it be edited and or deleted? I have been taught and can find no source contradicting the fact that pH goes down to 0. —Preceding unsigned comment added by 68.33.178.97 (talk) 04:55, 20 January 2009 (UTC)
Possibly better pH chart here: pH chart I created it and am very open to suggestions for change. --Hansepoo (talk) 02:57, 8 July 2009 (UTC)

Concentration < 10^-7

In Chemistry class today, when we were talking about the pH scale, I thought of a problem: If the concentration of H⁺ or OH^- is less than 10^-7M, taking the logarithm will result in it being over 7, which does not make sense, since how could an acid be diluted so much that is becomes a strong base, or vice-versa? I didn't see anything about it here, nor in the Chemistry book I looked at to find an answer, anyone have an explanation? --Gimmethegepgun (talk) 16:47, 11 March 2009 (UTC)

I think you need to focus on what dilution means. If you dilute an acid (pH < 7), you're adding water with a pH of 7, and making the resulting solution more like water than acid. Taken to the extreme, you'll basically just have water with a pH of 7. So dilution alone cannot increase the resulting pH beyond 7. Does that make sense? --PLUMBAGO 16:57, 11 March 2009 (UTC)

I understand that, and that's what doesn't make sense with the equation. The concentration will get smaller and smaller, but with a logarithmic scale, that would mean that eventually, according to the equation, the pH would be over 7, which isn't going to happen --Gimmethegepgun (talk) 17:02, 11 March 2009 (UTC)

The concentration of [H+] does not depend on the concentration of the acid alone. Write all your equilibria, ion, and mass balance equations before you solve for [H+]. --Cubbi (talk) 15:19, 1 May 2009 (UTC)

The fractional concentration of the acid's ions in the diluted solution will get smaller, yes, but they will be doing so against a background of the solvent, water, which has a pH of 7. So if you keep diluting something with water (homeopathy springs to mind), your solution's pH will tend towards 7. I think that you're actually asking a rather hypothetical question about what would happen if you could dilute an acidic solution with something that had no pH at all, such that the relative concentration of protons dropped below 10e-7. I think that this sort of solution breaks the assumptions of how pH is formulated, and that this "explains" the paradox. Unfortunately, what with me not being a chemist, that's a very poor articulation of your problem's solution. Can someone lend a hand?  ;-) Best regards, --PLUMBAGO 17:27, 11 March 2009 (UTC)

Yes, in nonaqueous solutions, the neutral pH is different from 7, if defined at all --Cubbi (talk) 15:19, 1 May 2009 (UTC)

Definitions

pH

pH is defined as minus the decimal logarithm of the hydrogen ion activity in an aqueous solution. By virtue of its logarithmic nature, pH is a dimensionless quantity.

$\mathrm {pH} =-\log _{10}a_{\mbox{H}}=\log _{10}{\frac {1}{a_{\mbox{H}}}}$

where a_H is the (dimensionless) activity of hydrogen ions. The reason for this definition is that a_H is a property of a single ion which can only be measured experimentally by means of an ion-selective electrode which responds, according to the Nernst equation, to hydrogen ion activity. pH is commonly measured by means of a combined glass electrode, which measures the potential difference, or electromotive force, E, between an electrode sensitive to the hydrogen ion activity and a reference electrode, such as a calomel electrode or a silver chloride electrode. The combined glass electrode ideally follows the Nernst equation:

$E=E^{0}+{\frac {RT}{nF}}\log _{e}(a_{\mbox{H}});\qquad \mathrm {pH} ={\frac {E^{0}-E}{2.303RT/F}}$

where E is a measured potential , E⁰ is the standard electrode potential, that is, the electode electrode potential for the standard state in which the activity is one. R is the gas constant, T is the temperature in Kelvin, F is the Faraday constant and n is the number of electrons transferred, one in this instance. The electrode potential, E, is proportional to the logarithm of the hydrogen ion activity.

--Mrfabioc (talk) 19:06, 3 May 2009 (UTC)

Applications: Carbon dioxide as an acid

The article states that carbon dioxide is an acid, but is it right to say this? My understanding is that carbon dioxide will dissolve in water to give an acidic solution of carbonic acid, but then carbonic acid is the acid, not carbon dioxide. Baclough (talk) 10:34, 24 May 2009 (UTC)

It is correct to say this under some of the definitions, but yes, the way it was phrased "water that has been exposed to air is mildly acidic. This is because water absorbs carbon dioxide from the air, and carbon dioxide is an acid. After absorption it is slowly converted into the weak acid, carbonic acid," was awkward. Rephrased to not say "acid" so many times. --Cubbi (talk) 18:54, 24 May 2009 (UTC)

Incorrect Information

The article says that, "The pH for pure water at 25 °C (77 °F) is close to 7.0.". I don't think that that is true. Pure water should always have a pH of 7, no matter the temperature. Please tell me if I am wrong.Nathan.tang (talk) 11:11, 13 July 2009 (UTC)

The pH of pure water depends on the equilibrium constant for the self-ionization of water. Equilibrium constants generally change with temperature and this one is no exception. --Itub (talk) 16:10, 24 August 2009 (UTC)