[teoporta]

Dear all,
I am using an organic monomer, and would like to have it with ph 3, 2 or 1 for example.
I'm not a chemist,  and would like to know if low ph could break the monomer...and how to fix such ph value.
Thank you!

p.s. any literature reference is very welcome

[Borek]

would like to know if low ph could break the monomer

Depends on the monomer.

and how to fix such ph value.

In general to keep pH constant you should use a buffer, but it won't be a trivial task at such low pH. For pH 1 you will probably get good results just by using 0.1M HCl (actually you will need slightly different concentration to account for the ionic strength). pH 2 & 3 buffers are difficult to calculate.

[magician4]

in addition:

for buffers with pH ~ 2 you might wish to take a look at hydrogensulfates:
HSO-₄^- / SO₄^2- 1:1 (mol/mol) should have a buffer point at ~ pH = 1,9
phosphorous acid / dihydrogenphosphate should show simmilair properties with pH ~ 2,19 at 1:1 (mol/mol)

for pH = 1 oxalic acid comes into mind, with a touch to the acid side in the buffer composition

pH=3 could be reached with HF/F- , though the extreme toxidity of this stuff from my point of view excludes it from any more serious consideration
i would prefer formic acid / formiate ~ 5:1 (mole / mole) for this task, accepting the somewhat weaker buffer properties and the asymmtric buffer behaviour over the toxidity of HF/F-

regards

Ingo

[Borek]

for buffers with pH ~ 2 you might wish to take a look at hydrogensulfates:
HSO-₄^- / SO₄^2- 1:1 (mol/mol) should have a buffer point at ~ pH = 1,9
phosphorous acid / dihydrogenphosphate should show simmilair properties with pH ~ 2,19 at 1:1 (mol/mol)

This is not wrong, but it is potentially misleading. Preparing solution that is 1:1 HSO₄^-/SO₄^2- is not as obvious as it looks (and not only because of the high ionic strength of the solution).

[magician4]

This is not wrong, but it is potentially misleading.

maybe I'm missing something here, but:

both sodium hydrogensulfate and disodium sulfate are commercially available products

mixing 0.1 mole of each, and adding water ad 1 L will result in the 1:1 situation (with respect to the c₀ 's)

due to this pH-calculator (which calculates a little deeper than just Henderson-Hasselbalch) :
http://de.webqc.org/phsolver.php

( values inserted were :
HSO4- pKa=1.9 c=0.1
SO42- pKb=12.1 c=0.1 )

... this should result in a pH of 1,989...

good enough for pH ~ 2 in my opinion

what do I miss?


regards

Ingo

[camptzak]

will the higher ionic strength of the solution effect the activity of the H⁺ ions? and if so, how? could you modify the Henderson-Hasselbalch equation factoring in ionic strength?

and for the person who stated the thread, could you show us the monomer you are planning on using?

[teoporta]

Thank you for your responses,

i'm using pentaerythritol etoxylate http://www.sigmaaldrich.com/catalog/product/aldrich/416150?lang=ja&region=JP

i'd like to obtain different pHs to observe nanoparticles' dispersion behaviour.
But it becomes useless if the monomer breaks.

[camptzak]

Thank you for your responses,

i'm using pentaerythritol etoxylate http://www.sigmaaldrich.com/catalog/product/aldrich/416150?lang=ja&region=JP

i'd like to obtain different pHs to observe nanoparticles' dispersion behaviour.
But it becomes useless if the monomer breaks.

It seems like the only thing you might run into at a low pH is an intermolecular reaction creating a cyclic ether an releasing water. I dont know how low the pH has to be in order   for this to happen.
what solvent are you going to be doing your experiment in?

[magician4]

will the higher ionic strength of the solution effect the activity of the H⁺ ions? and if so, how? could you modify the Henderson-Hasselbalch equation factoring in ionic strength?

of cause ionic strength is a factor always to be considered, though with solutions at 0.1 mol/L more often than not we tend to keep on working with concentrations instead of activities.
there is a longstanding debate exactly when the difference between activity "a" and concentration "c" becomes a relevant factor for respective calculations, but for those who love more exact calculations, Debeye-Hückel will give you instruments for a first refinement.

after that, the activity (instead of the concentration) can be introduced in the usual terms for pH-calculations

I would NOT advise to use the standard Hendersson-Hasselbalch equation for calculations regarding sodium hydrogensulfate / di sodiumsulfate, as esp. the degree of dissociation of KHSO₄ can't be neglected

I would propose to use a full law of mass action term for the calculation instead (see attachment) and solve the resulting equation


regards

Ingo

[teoporta]

was thinking about 2-butoxyethanol as solvent...

[Borek]

both sodium hydrogensulfate and disodium sulfate are commercially available products

mixing 0.1 mole of each, and adding water ad 1 L will result in the 1:1 situation (with respect to the c₀ 's)

due to this pH-calculator (which calculates a little deeper than just Henderson-Hasselbalch) :
http://de.webqc.org/phsolver.php

( values inserted were :
HSO4- pKa=1.9 c=0.1
SO42- pKb=12.1 c=0.1 )

... this should result in a pH of 1,989...

good enough for pH ~ 2 in my opinion

Simple approach, ignoring ionic strength (but see below):

       HSO₄^-        H⁺      +     SO₄^2-

I:       0.1               0              0.1
C:       -x                x               x
E:     0.1-x              x            0.1+x

[tex]K_a = \frac {[H^+][SO_4^{2-}]}{[HSO_4^-]} = \frac {x(x+0.1)}{0.1-x} = 0.01[/tex]

Solving for x we get 0.00844, so pH is 2.07 - and I agree for most cases this is close enough to 2. But if you will try the same approach for 0.01 M solutions, pH will got up to 2.38, so we are almost 0.4 pH unit away, even if 0.02M buffer looks quite reasonable. HSO₄^- is a relatively strong acid, and it makes calculations difficult.

Note: I got the same pH using full blown equilibrium calculations using pH calculator built into Buffer Maker. I am not going to comment on the phsolver result - but the idea of entering Ka and Kb looks absurd for me, as they describe the same equilibrium (so it is enough to enter one of these numbers).

Now, ionic strength. For 0.1M solutions Buffer Maker calculated ionic strength to be 0.48 - well above the limits for Debye-Huckel theory. BM uses Davies equation, which is sometimes good for up to 0.5 (although some researchers say it is good to 0.3 only). Assuming it is good for ionic strength of 0.5, pH of the 0.1/0.1 solution is around 1.74 - whether it is still a good approximation of pH 2.0 is debatable.

For 0.01/0.01 solution ionic strength is calculated as 0.052, so we are in the Debye-Huckel theory range (up to 0.1), but pH calculated is 2.32, so again it is not clear if it is an OK approximation of pH 2.0.

It is perfectly doable to get pH=2.00 solution using SO₄^2-/HSO₄^- - just use 0.024M hydrogensulfate and 0.01M sulfate. As I wrote earlier - preparing such a solution is perfectly possible, but not as obvious as it may look.

I would NOT advise to use the standard Hendersson-Hasselbalch equation for calculations regarding sodium hydrogensulfate / di sodiumsulfate, as esp. the degree of dissociation of KHSO₄ can't be neglected

But that is - basically - what you did, assuming 1:1 solution will yield pH close to pKa.

I would propose to use a full law of mass action term for the calculation instead (see attachment) and solve the resulting equation

Agreed. That's what I did (even if for pedagogical purposes I posted ICE method here, but I checked it yields the same results).

Please see this thread - LaTeX can be entered directly into posts.

[Borek]

was thinking about 2-butoxyethanol as solvent...

That makes things even more complicated - pH is defined for water solutions, there is no such thing as pH of non-water solution.

[magician4]

Simple approach (...)
[tex]K_a = \frac {[H^+][SO_4^{2-}]}{[HSO_4^-]} = \frac {x(x+0.1)}{0.1-x} = 0.01[/tex]

I take it that we both agree, that we both did put up the same expression for calculating this problem, except you did use 0.01 instead of 10^-1.9  for the K-value (which is quite in the same ballpark, I agree)

Solving for x we get 0.00844, so pH is 2.07 - and I agree for most cases this is close enough to 2. But if you will try the same approach for 0.01 M solutions, pH will got up to 2.38, so we are almost 0.4 pH unit away, even if 0.02M buffer looks quite reasonable. HSO₄^- is a relatively strong acid, and it makes calculations difficult.

I would look at the "difficulty" from another point of view: that's just the way stronger acids (compared to for example acetic acid) do behave - and that's why they're giving acceptable buffers only with high concentrations (and that's why I was recommending 0.1 mol/L in the first place).
the reason is that they do dissociate in a relevant manner - hence c ~ c₀ doesn't apply anymore and the situation becomes concentration - dependant ( and not only ratio acid : corresponding anion dependant)
on the other hand, the very calculation itself doesn't seem difficult to me

anyway, talking pH 2 or - even worse - pH 1 , these acids are the only game in town, so we'll have to roll with the punch

Note: I got the same pH using full blown equilibrium calculations using pH calculator built into Buffer Maker. I am not going to comment on the phsolver result - but the idea of entering Ka and Kb looks absurd for me, as they describe the same equilibrium (so it is enough to enter one of these numbers).

that's just the way you have to introduce a buffer in this calculator: entering the c₀ - values of the acid, the corresponding anion and their respective K values

Now, ionic strength. (...)

My personal experience with buffers is, that you'd better calculate a value of expectation first, and then start preparing the system by taking one component's solution (usually the corresponding anion) and add the second component under pH-control.
carbon dioxide, the real temperature, other impurities, quality/purity of the educts...
the list goes on and on
no use to calculate the theoretical pH to the third decimal  with D.H. or other approaches (except for academic reasons, that is) , if those effects will spoil everything in the first decimal already.

... and my impression was, that teoporta was asking for practical help

I would NOT advise to use the standard Hendersson-Hasselbalch equation for calculations regarding sodium hydrogensulfate / di sodiumsulfate, as esp. the degree of dissociation of KHSO₄ can't be neglected

But that is - basically - what you did, assuming 1:1 solution will yield pH close to pKa.

I didn't need a HH-calculation for the estimation of the buffer point, knowing that the rule of thumb "an acid will buffer at / near it's pKa" will be +/- valid for acids of medium strength in higher concentrations

I would propose to use a full law of mass action term for the calculation instead (see attachment) and solve the resulting equation

Agreed. That's what I did (even if for pedagogical purposes I posted ICE method here, but I checked it yields the same results).

Please see this thread - LaTeX can be entered directly into posts.

So, by the end of the day we both come down to the same recommendation, both with respect how to calculate this system and what needs to be done - except that i would recommend higher buffer strength than you do
but that's no basic disagreement, I take it

... and that is good for teoporta


regards


Ingo


p.s.: thank you for directing me to a thread explaining how to use LaTex in this forum: that's what I've been looking for

[magician4]

was thinking about 2-butoxyethanol as solvent...

you could test for any adverse effects by simply testing this solvent "stand alone" for the respective "pH-values" and reaction conditions (as Borek already pointed out: definition of pH relates to waterbased systems exclusively, though there are alternative systems for other solvents), as the terminal -O-CH₂-CH₂-OH group of butoxyethanol should behave quite similar to your polyethoxylates

pls. also note that pKa values from waterbased systems can not be used for other solvents, hence any "buffer" might need a completely different composition from those already mentioned here


a final remark:

It seems like the only thing you might run into at a low pH is an intermolecular reaction creating a cyclic ether (...)

I would expect polymerization to a duroplastic network instead, as with large rings it is extremely unlikely to have their formation win over such a process - or, more likely , reaction with your solvent to form butyl-endcaped polyethoxylates

... if anything bad happens at all: primary alcohols are not the most sensitive guys around if you don't heat 'em too much, even under acid conditions


good luck

Ingo

[Borek]

I take it that we both agree, that we both did put up the same expression for calculating this problem, except you did use 0.01 instead of 10^-1.9  for the K-value (which is quite in the same ballpark, I agree)

My sources give pK_a2 as 1.99.

I would look at the "difficulty" from another point of view: that's just the way stronger acids (compared to for example acetic acid) do behave

Yes, problem is, people often forget about it, and end with absurd results.

anyway, talking pH 2 or - even worse - pH 1 , these acids are the only game in town, so we'll have to roll with the punch

I tend to disagree. At low (and high) pH you don't need buffer to keep pH constant. Buffer capacity of these solutions is high enough on its own:



That's a buffer capacity curve for 0.1M acetic buffer - note, that at pH around 1.6 buffer capacity of the solution becomes comparable with buffer capacity at pH 4.75, even if at 1.6 acetic buffer is no longer working; just the presence of H⁺ in relatively high concentration makes the solution resistant to pH changes. You don't need an acid/conjugate base pair.