During optimization of an extraction procedure, we found out that salt concentration has a significant effect on the pH value of buffer solutions. Although this effect is simply explained in some sites, after intensive research I wasn't able to find any reliable manuscript which clearly explains the effect. Can anyone refer me to one?
As far as I'm concerned, the Salt in the Henderson Eq refers to the conjugate base or acid concentration in the buffer solution, I have no idea if the NaCl concentration plays any role in the buffer equilibrium, and i don't think we can explain it by co ion effect, this is what i exactly found:
"It depends upon the pH of the water into which the salt is being introduced. Chemically speaking, salt is a base compound, falling smack in the center of the acid-alkaline spectrum. If introduced to water which has a high pH, the pH might be lowered incrementally toward the center of the pH spectrum depending on how much water there was and how much salt was introduced. If the water had a very low pH, making it very acidic, the salt would increase the pH toward the center of the spectrum" but I can't find any source for it
.. addition of salt of strong acid and strong base like NaCl to water does nothing to pH becoz the species formed OH- and H+ balance each other out . however if salt of strong base and weak acid is added to water like ( Sodium acetate) then addition of this salt will result in basic pH.
and if salt is of weak acid and strong base like (ammonium chloride) is added to water it will result in acidic solution.
Refer to hydrolysis of salts in any Physical Chemistry textbook....
I too am experiencing false high pH in high saline water.
We deal with Acid Sulfate Soil issues affecting water.
I commonly see phenolphthalein change colour at pH 9.4 in high saline with high acidity water environments. The question then becomes is the pH true because we all know phenolphthalein changes colour at pH 8.3?.
Be very interested if you pursued your question and got a definitive answer?
It all can be explained by Debye-Huckel eq and concerning the activity of ions rather that their concentration, It is the exact word which I've written in my paper and i think they are right,
"This interaction can be explained by the effect of increasing the salt concentration on the pH of aqueous solutions. In plain words, by raising the ionic strength of an aqueous solution, based on the Debye– Huckel equation, the activity of the hydroxide ion decreases. In other words, by adding salt to an aqueous solution with a high or low pH due to increasing the ionic strength, the pH drifts towards the center of the pH spectrum (pH= 7)"
It 's also match with your experience,I suggest adjusting the pH after the adjusting the salt concentration in that way the nominal pH might be equal to actual one. The phenolphthalein color changes at pH 8.3 and nothings changes it so you can deduce that the actual pH is different from what you thought.
Good day everyone,
I stumbled upon this article, but it was published back in 1929. Perhaps this might be relevant? Hope this helps
thanks
Hi, It was long time ago and I have already changed my field, but what i have seen was an interaction which DOE has shown to us, we needed an explanation , I mean why something like this has happened? I performed the experience again and also check some points out of design,, the results was clear if you set the solution pH and then add some amount of salt (i mean neutral salt) the nominal pH drifts toward 7, The Deby-Huckel was a good explanation, it states by increasing the ionic strength, the activity will decrease (obviously) and what pH meter shows and actually what really happen in solution is all about activity not concentration, Since then I just recommend that if you want to have the accurate pH, set at the end of the experiment, or check it after adding the any other substance.
Li Shu,
Are you able to please contact me via email [email protected]
I commented back in April 11 2014 on this subject and note you are from RMIT.
Cheers
Paul
Hi, LI Shu, I attached my article, all details are included, although it is an analytical article and therefore the final purpose is different, but i think it leads to good results, I mean maybe it was my experimental mistake, I had to check the pH at final step (I trusted the buffer solutions), but the results are clear you can see the interaction in response surface plot and as it is clear in the ANOVA table they interaction is consider as significant, also it should be mentioned that duo to the nature of the analyte, pH has significant effect on it's behavior, but also salt concentration plays very important rule, if something is unclear, just ask if i can help i will be very happy, bests
Just wanted to chip in an alternative explanation for those experiencing similar problems in the future. When using strong acids/bases to adjust pH such as HCl, the dissociation % should be near 100% regardless of how much salt is added to the solution. Here I cite from wikipedia,
"When chloride salts such as NaCl are added to aqueous HCl they have practically no effect on pH, indicating that Cl− is an exceedingly weak conjugate base and that HCl is fully dissociated in aqueous solution."
As Ghazaleh Aliakbarzadeh mentioned, increased activity of protons at high ionic strength may produce artifacts in pH measurements where the measured pH is not exactly the concentration of protons in solution. But another situation to watch for is when pH is measured in solutions of very low ionic strength, where the millivolt output from proton gradients are too small to be picked up by pH probes. For example, if you measure a 0.1M HCl solution in distilled water, you will get readings approx. 0.05-0.1 pH units higher than the actual value, let's say pH 1.08. As you add NaCl, you may observe the pH decreasing to pH 1, but in reality, it is only approaching the real pH as the milivolt output signals increase, within the sensitivity range of the pH meter. Of course, this is only the case for strong acids. I do believe salt concentration will actually change pH when weak acid/bases are involved, as Krishnasarma pathy explained.
Hi Ghazaleh! I just can't believe that I have the same problem as yours in an emulsion system, the pH is affected significantly by the salt and salt concentration, so, did you finally accept the theory that salts are bases and they tend to pull the pH to the neutral 7? Did you had any problem backing that up with that theory? THANK YOU!
Hi Wilfredo! As I mentioned in my previous replies we concluded that salts are neutralizers although it may just be the case for neutral salts like NaCl not all of them, they drift the pH toward 7, I am pretty sure about that and I can just recommend you to adjust or check the pH of your samples at final step prior to analyse or any other consideration even if you have buffer solution, Good luck
Hi Miss Aliakbarzadeh,
thank you for your describe about your experiment.
can you say to me which following salts has more effect on the pH of solution?
NaCl or KCl
if you have good article about this title. it's my pleasure to use it.
Regards,
jafari
Dear Ghazaleh Aliakbarzadeh,
Our paper titled "Why Does pH Increase with CaCl2 as Draw Solution during Forward Osmosis Filtration" (doi:10.1016/j.psep.2016.06.007) has just been online(http://authors.elsevier.com/sd/article/S0957582016300933).
The pH of neutral salts is considered as 7. In the paper we showed that the pHs of both NaCl and Na2CO3 increased with concentration while the pH of CaCl2 decreased with concentration. We proposed that it was possible for all substances to have their unique pH values.
Hope it helps.
Li Shu,
Read with interest your recent publication and in particular your comment:
"" In contrast to conventional belief, the standard curves showed that pH of NaCl increased with the increase in concentration""
This was my observation in April 201 discussion as above.
I am thinking it is competition with H+ by other monovalent cations and in this case excessive Na+ ions.
Wont happen with multivalent cations.
What are your thoughts?
Cheers
Paul
Li,
This is intriguing and I have still yet to find anywhere an explanation of the pH of NaCl increasing with the increase in NaCl concentration. Yet your observation is the same that I got in the field 3 years ago. Being a scientist i would like to know why, but as you have pointed out, the observation is contrary to conventional belief and understanding.
Going back to basics, the difference in the potentials of the pH and reference electrodes provides a millivolt signal proportional to pH. It is only H+ that interacts at the pH (glass) interface to produce a potential (voltage) not Na+ or other multivalent cations. We must exclude ammonia, bromide,iodide and sulphides ions for their known deleterious effects on the reference electrode at this stage.
We also know that sodium error occurs at very high pH where the hydrogen ion concentration is very low in comparison to sodium ion concentration.The sodium ion concentration can be so high relative to hydrogen ion contration that the electrode begins to respond to the sodium ion. This results in a reading that is lower than the actual pH and is the reverse of our observations, (ie behaves as per conventional belief).Under a pH of about 10.0 sodium error is not expected to happen.Expect sodium error usually in extreme pH ranges eg. >pH13.
In high salt concentrations, are there sufficient Na+ ions present to competitively restrict H+ ions from reaching the glass interface????
I'm open to any input on this,...... and havent as yet been able to find any literature evidence on why this observation occurs.
Thank you for your time once again
The following two reports are worth your consideration, even if only as the starting points for a forward literature search. Will post anything else I come across, to help suss whether the effect is instrumental or physicochemical in each of your cases.
http://pubs.acs.org/doi/abs/10.1021/ac50164a034?journalCode=ancham [F. E. Critchfield, J. B. Johnson, Effect of Neutral Salts on pH of Acid Solutions, Anal. Chem., 1959, 31 (4), pp 570–572; DOI: 10.1021/ac50164a034]
http://www.jbc.org/content/82/3/775.full.pdf+html?sid=2a02e597-ad28-4855-81e2-cbbd36feafe5 [H. W. Robinson, The Influence of Neutral Salts on pH of Phosphate Buffer Mixtures, J. Biol. Chem., 1929, 82:775–802; also suggested above by SV Zulfigar]
Mechanism/explanatory papers harder to come by; perhaps you will find them before I. Note, affiliation with NU is now former. Cheers. MED
Post Script: If I recall, the activity series observed in at least one of these was Li > Na > K, consistent with some of the comments above. Cheers. MED
MED,
Thanks for your contribution, i really appreciate your input on this observation.
I read with interest the two journal articles and they support the conventional belief and understanding where the pH decreases.
Still no evidence cited for an increasing pH.
Cheers
Paul
By definition pH is based on activity of a solution an not the concentration of H+.
At 25 degrees and 1 atm:
Kw={H+}{OH-}/{H2O} and {H+}={OH-}
{H+}=(Kw)^0.5
pH= -0.5*logKw= 7.00
At colder temps (e.g. 4 degrees), the van T'off equation applies to the relationship between temperature and reaction rate at an assumed constant enthalpy.
pH=7.37
At increased ionic strength (usually < 0.1M; e.g. NaCl) activity can be assumed to be ~1 still, so the pH shouldn't change and = 7.00
At higher ionic strength, the Debye-Huckel Equation applies, and activity (gamma) changes:
log gamma= -Az2*I0.5; where A=0.5 at 25 degrees, z= charge, I= ionic strength
So in sea water where I= 0.5-0.7M, the pH will increase because activity increases:
So pH= 8.08 to 8.33 on various scales depending on the inclusion of various anionic hydrogen complexes (e.g. HSO4-)
Alistair,
Thank you very much, you have nailed it.
I'm a biochemist with medical background and couldn't find in the literature an explanation with a scientific authentication. On reading more (background) on the information you have supplied the gamma activity does explain the increase in pH as ionic strength increases.
I will sleep well tonight!
Thank you to all that have contributed to this vexing problem I observed in the field!
Hi Paul,
I'm an environmental analytical chemist who teaches Biochemistry. All I could hear was my aquatic Chemistry professor yelling at all of us for not saying "activity" right off the bat. 😉 I'm glad it was helpful
Yes, they are the teachers we will always remember...
All the very best.
Is right to say following::
With increase in salt concentration in the solution ionic strength increases while keeping pH constant?
I believe it's not ok to say that unless you justify your answer. As outlined above. Within a certain range of ionic strength
Hi Alistair,
When you say:
"So in sea water where I= 0.5-0.7M, the pH will increase because activity increases"
shouldn't that be - "ph will increase because activity decreases"?
According to the DH formula as ionic strength (I) increases won't gamma (activity coefficient) decrease?
gamma = 1/e(Az^2*I^0.5)
Then the activity of the H3O ions decreases leading to an "increase" in pH registered by the pH probe.
Or am I missing something?
I believe if the way I'm reading your question is correct, the equation isn't the inverse.
It's just Gamma= e(-Az^2* I^0.5).
Hope this helps. So the previous comment stands.
I am doing experiment with the buffer solution at pH around 10.2. Addition of 30 g/L NaCl to the buffer solution resulted in a decrease in pH values to about 9.8. However, based on the calculations Alistair has done using Debye-Huckel equation, the pH value actually should rise by addition of salt.
Also, based on what Paul said "We also know that sodium error occurs at very high pH where the hydrogen ion concentration is very low in comparison to sodium ion concentration.The sodium ion concentration can be so high relative to hydrogen ion contration that the electrode begins to respond to the sodium ion. This results in a reading that is lower than the actual pH and is the reverse of our observations, (ie behaves as per conventional belief).Under a pH of about 10.0 sodium error is not expected to happen.Expect sodium error usually in extreme pH ranges eg. >pH13"
So, is there any other equation to be used for high pH or is the pH drop as a result of instrumental error?
The pH of a buffer depends not on the concentration of the buffer ions but on a thermodynamic quantity called the activity coefficient. This parameter is strongly affected by the total concentrations of ions in the solution. Some general-purpose glass electrodes are almost always somewhat permeable to sodium ions. Therefore, a potential related to the Na+ concentration can be produced in the same way as with H+ ions- If Na+ is present in a solution whose pH is to be determined, the measured pH decreases as the N+ concentration increases, because the electrode detecs the sum of the H+ and Na+ concentration.
We need to understand the following points:
1.
The pH meter measures activity of [H+].
2.
In diluted aqueous solution [H+] concentration equals to its activity, ie, the activity coefficient of [H+] close to 1. In such diluted solution, pH reflects solution [H+] concentration therefore how acidic it is. In the meantime, the diluted solution has the relation of [H+] * [OH-] = 10^(-14). This solution has high pH means it has low [H+] and results in high [OH-] so we say this solution is basic.
3.
In concentrated aqueous solution [H+] concentration is higher than its activity, ie, the activity coefficient of [H+] smaller than 1. In such diluted solution, pH reflects solution [H+] activity. The high pH measured by pH meter indicate low [H+] activity but not high [OH-] activity. This solution has high pH does not mean high [OH-] concentration so we CANNOT say this solution is basic. Imagine the pH meter measures [OH-] activity, it will show low as well.
4.
If we really want to know the pH of the very high concentrated solution, the easiest way is to dilute the solution to give a meaningful pH.
On the effective buffer pH range for weak acid / strong base buffer solution: https://www.researchgate.net/post/why_is_making_buffer_at_pH_pKa1_is_recommended
On deriving the Henderson-Hasselbalch equation equation, along with a justification for the minimum concentration limit for its validity: https://www.researchgate.net/post/pH_calculation_of_a_mixture_of_formic_acid_NaOH_and_water
On the pH of a solution of trisodium phosphate (Na3PO4) and NaOH, possibly obtained by titration of phosphoric acid with NaOH: https://www.researchgate.net/post/Im_looking_for_a_good_pH_13_buffer_recipe_for_hydrolysis_experiment
About predicting the pH of citric acid ― sodium citrate solutions and buffers:
(i) for approx. pH < 4.0 ― cf. my post at:
https://www.researchgate.net/post/How-to-prepare-a-citric-acid-solution-of-pH-25-and-3-from-100-citric-acid-powder
(ii) for approx. 4.0 < pH < 5.5 ― cf. my posts at:
https://www.researchgate.net/post/Why-the-pH-of-Citrate-buffer-increases-when-diluted-with-water
(iii) for approx. 5.5 < pH < 8.0 ― cf. my posts at:
https://www.researchgate.net/post/Why-the-pH-of-Citrate-buffer-increases-when-diluted-with-water
(iv) for approx. pH > 8.0 ― cf. my posts at:
https://www.researchgate.net/post/Why-the-pH-of-Citrate-buffer-increases-when-diluted-with-water
(v) for approx. pH > 9.0 ― cf. my posts at:
https://www.researchgate.net/post/Why-the-pH-of-Citrate-buffer-increases-when-diluted-with-water
I made some Ph measures of different NaCl salt solutions 1%, 2%, 3 %, 4%, 5% and 36%. Surprisingly I got very different results from 6,0 (1%) to 8,7 (36% NaCl solution). I got the 7.1 Ph by the 3% NaCl solution. My opinon: it is based a little bit of the sensitiveness of the electrode. For example if the electrode is not sensitive enough, a low concentration of the solution NaCl like 1% will not give the exact outcome of 7,00 to 7,4. If You are measuring only deionizised water You will need a special electrode. Otherwise the result will be more than surprising (very far from 7,0). We all know that the solutions are made with deionizised water (the H+ are very weak)
https://www.researchgate.net/post/Does_NaCl_change_pH_Does_NaCl_effect_pH_Does_NaCl_salt_change_the_pH_in_a_buffer_solution
I think there is an issue that may not have been considered in depth: is the pH probe measuring the real pH or not? H+ and Na+ (or K+ or any other monoatomic cation M+) will compete with each other to cross the glass membrane and produce a response at the pH probe junction. If the concentration of Na+ (or any M+) is much larger than the concentration of H+ in solution, much of the pH probe response may come from the M+ and not the H+, which would explain the correlation between larger M+ concentration and larger apparent pH measurement. In this case, the problem is inherently due to the design of most pH probes: they are not specifically selective to H+ only, but to any M+ that produces a voltage response in the pH meter. If a buffer is involved in the solution, everything can change if common ion effect is possible.
Alejandro Cuello Yes! Encountered this myself recently.
Article Quantifying the Cross-Sensitivity of Glass pH Electrodes in ...
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