Discussion about pH

So what is the question you seek to be answered?

You stated a zero reference point or a base line in the mesearment of pH, being "14".

I would believe this is an actual reference of natural standing water, possibly being an average found around the globe. So you figure, some people did a whole great deal of research to actually find this to being the known average of pH saturation.

So, in essence, this has been excepted as the "zero base-line" in the measurement of pH, and that is that.

Maximo

There are very specific changes between a fluid being alkaline and acidic.

Right in the middle of this change over is neutral with a pH of 7.0, a fluid can not be more acidic than 14 or 0.

I forget which way round the scale is now... Ooops!

John.

The pH value is defined as pH=-log[H+] where the log is base 10 and [H+] represents the concentration of hydrogen ion in gm mole/liter. The dissociation of water at 25 deg Celsius results in an equal amount of hydrogen ions and hydroxyl or OH- ions.

That given, it is found that at neutral condition for absolutely pure water at 25 deg C, there are concentrations of [H+] = [OH-] = 1.0 x 10-7 gm mole/liter of these ions. I do not know of any reason why the concentration of hydrogen ions could not decrease to less than 1.0 x 10-14 , but it would be exceedingly unlikely. On the other hand, a pH=0.0 suggests the ability of an infinitely high concentration of hydrogen ions since the math would require divison by 0.0, therefore we have the condition of the unattainability of zero pH.

In reply to Aqua Doc: No division by zero.

When pH =0, the activity of hydrogen ions is 1 (lg(1) =0) !

What happens when the activity of H+ is 1.2 ? Or 2 ? Or ...

The only one fact that is well stated is that the product of the activity of H+ and HO- , [H+]*[HO-] = constant at a given temperature (constant = 10^-14 at 25 deg Celsium) and by definition pH = -lg[H+].

The question is: can the pH of a aqueous solution be less than 0 or higher than 14 ?

For example: for HCl aqueous soln of 1 mol/l = 1 Eg/l (aprox 3.85%w), pH will be

pH = -lg(1) = 0. What if the concentration will be aprox 1.5 mol/l = 1.5 Eg/l (aprox 6%w) ?

Math: pH=-lg(1.5)=-0.17.

Some problems (theoretical and practical) will rise due to the difference between concentration and activity, but let's try ignore them here or condensing making the problem more specific:

1.For an acidic aqueous soln can be the activity of H+ greater than 1 ? 1.5 ?

2.If yes: pH = ? I don't know the answer.

Remark: In school:

"pH is between 0 and 14". OK.

"...and is -lg[H+]". OK, but we have a problem.

"a pH=0.0 suggests the ability of an infinitely high concentration of hydrogen ions since the math would require divison by 0.0, therefore we have the condition of the unattainability of zero pH"

Please re-check.

A pH of 0 simply means that [H+] = 1gm mol/litre. There are 2000/18gm of hydrogen in a litre of pure water, assuming an SG of 1.000. So, even at pH=0, there is still a large amount of un-dissociated hydrogen present.

OOps! Sorry, My Bad.

I remember to agua doc, ICPEAR, PWSlack, et cetera:

The pH is NOT = log(H+) !!!

The argument of log (exp, sin, etc) MUST BE A-DIMENTIONAL !!!

So, the correct definition of pH is: log (con-H+/conc. of water=1).

The use of logarithm avoids the use of numbers in exponential notation. It is the same, when you treat the RATIO (A-dimensional !) of electrical, sonic etc. quantities using deciBel.

In principle, pH < 0 or > 14 is possible.

Agreed.

I believe he is asking about seeing pH levels quoted outside the acid-base range. Is he misreading? Or accurately reading inaccurately typed/edited material?

pH is a measure of the acidity or basicity of a solution. Solutions with a pH less than 7 are considered acidic, while those with a pH greater than seven are considered basic.

pH 7 is defined as neutral because it is the pH of pure water at 25 °C this mean the quantity of OH^- and H⁺ ions are equal.

pH is formally dependent upon the activity of hydrogen ions (H⁺), but for very pure dilute solutions, the molarity may be used as a substitute with some sacrifice of accuracy.

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.

The pH reading of a solution is usually obtained by comparing unknown solutions to those of known pH, in the case of an electrode the diferential of potential between the reference electrolite and the sample is the measure of pH value, If you have a sample with more than 14 pH units then you have an alkaline solution and the quantity of OH^- ions may be more than 14 units but still be an alkali, the same for an acid solution.

The concept of pH was first introduced by Danish chemist Sorensen. However pH is actually a shorthand for its mathematical definition, in chemistry a small p is used in place of writing − log₁₀ and the H should more correctly be [H⁺], standing for concentration of hydrogen ions.

My experience with ph is in home, shop and garden applications. deionized water has a ph of about 5.2. Tap water has a ph of 7 here. The cyanide compounds I use for electro-plating are all about a ph of 14 and balsamic vinegar comes in at a ph of 4. My garden soil has a ph of 7.2 so I water my garden with water that is ph adjusted to 6.5 for most of my garden. I get the lower ph by deionizing some of the water provided through drip lines. I was not surprised that rain water and snow have a ph of 5.2-5.4. I find it curious that HCL would have a PH of 0-1 and water a ph of 5.2 when oxygen and chlorine are so close chemically excepting the valences of chlorine is 1,3,5,& 7 while water has a valence of 2. Since ph is a measure of hydrogen ions, how is this difference explained?

In reply to Aqua Doc: No division by zero.

When pH =0, the activity of hydrogen ions is 1 (lg(1) =0) !

What happens when the activity of H+ is 1.2 ? Or 2 ? Or ...

The only one fact that is well stated is that the product of the activity of H+ and HO- , [H+]*[HO-] = constant at a given temperature (constant = 10^-14 at 25 deg Celsium) and by definition pH = -lg[H+].

The question is: can the pH of a aqueous solution be less than 0 or higher than 14 ?

For example: for HCl aqueous soln of 1 mol/l = 1 Eg/l (aprox 3.85%w), pH will be

pH = -lg(1) = 0. What if the concentration will be aprox 1.5 mol/l = 1.5 Eg/l (aprox 6%w) ?

Math: pH=-lg(1.5)=-0.17.

Some problems (theoretical and practical) will rise due to the difference between concentration and activity, but let's try ignore them here or condensing making the problem more specific:

1.For an acidic aqueous soln can be the activity of H+ greater than 1 ? 1.5 ?

2.If yes: pH = ? I don't know the answer.

Remark: In school:

"pH is between 0 and 14". OK.

"...and is -lg[H+]". OK, but we have a problem.

pH 0 to 14 is valid for H2O as this is directly associated with the dissociation equilibrium of H2O (molecular) into its ions H+ and OH-

If you multiply the concentration of both ions, it will be 10 exp(-14) at standard temperature. Therefore, it is not possible to have more than pH 14 in pure water at standard temperature.

This is not the same for other substances and it is usual to get pH values of 20 or more for some organic solutions.

This is the reason for some instrument manufacturers to have instruments with specs showing wider pH ranges.

If anyone likes to read more on this, I would advise to read the white papers from many of the pH meter manufacturers or even better, books on this subject.

I must insist: to understand pH measurements, it is absolutely necessary to understand the dissociacion equilibrium of the molecule into H+ ions.

I agree that things have to change for organics/ other solvents that water.

But: the ionic product for water is 10-¹⁴ but there are no visible restriction for H+ or HO- concentrations derived from that, except their dependace one of the other. Can you see any ? Why do you establish a link between pH and the ionic product ? pH =-lg[H+] YES; pH= -lg[H+][HO-] NO.

[H+]*[HO-] = 10^-14, water,25 deg Cesium. Right ?

[1.25]*[8^-15] = 10^-14. Right ? Water is water, we agree here. See below:

pH = -lg(1.25)= - 0.097 .

If pH was -1, then pOH would be 15.

If pOH were -1 then pH would be 15.

Simply because [H+][OH-] = 10^-14. This is the Dissociation Product of water.

Does this make sense?

pH +pOH = 14 , water, at 25 deg C accordingly ti diss product [H+][OH-] = 10^-14.

So, if we assume pH=-1 -> pOH=15 it's correct. Also, mathematically pH=15 -> pOH=-1 is correct from the point of view of the water diss product.

The restriction you have is that pH depends on the dissociacion equilibrium of the solvent (for instance, water)

When you displace the dissociation equilibrium to the acidic region, you will have a higher concentration of H+ ion and a lower concentration of the OH- ion, but always the ionic product is the same (temperature considered) (10 exp(-14), and the pH range remains 0 - 14

pH is another way to show H+ ion concentration but as this is related to the dissociation of the molecule (H2O in our example), we have to respect the range given by the ionic product. In the same way, we can address the pH range for any other solvent that has a dissociation equilibrium with its H+ ion. In other words, we should know its ionic product which is usually different from 10 exp(-14)

That is why the pH range is solvent dependent.

A standard Ph meter is usually calibrated from 0 to 14 units.The denotes the power(exponent) of ten of the Hydrogen or Hydroxl ions. 7 is neutral, 0 is very acidic, and 14 is very caustic.(lye).A standard Ph probe produces 59.33 mv per Ph unit, and substitution boxes are manufactured to this value to troubleshoot probes and meters.The last 2 digits to the right of the decimal point may be off a little, since I am working from memory, and it has been a long time since I calibrated a lab Ph meter.

The point is that most meters adhere to the 0-14 standard.