Chromatography Forum

Hi,

we are using a analysis method for a charged molecule (ammonium, Cl-) with ACN and TFA buffer (0.1% or 0.05%).
We would like to make a Prep LC on this product to separate impurities without changing counter ions (Cl-), do you think we can replace TFA buffer with HCl buffer in the same percent ?
Thanks.

Neither 0.1% TFA nor 0.01M HCl are a buffer. However they both have sufficient concentration of H+ that they do maintain the pH at around 1. There are however two practical problems for you. First, TFA is an ion-pairing acid and HCl is not, therefore the selectivity will be different, and likely the peak shape worse. You will need to work out new conditions. The more annoying problem is that HCl is much more corrosive to stainless steel than TFA. You will want to use titanium, PEEK, teflon and glass for the wetted surfaces. (For a one-time job, you probably can get away with a passivated 316-stainless system.)

Provetech,

What is your column? if you are working with ion exchange or mixed mode column you can replace TFA with any other acid at different concentration (expect some change in selectivity and retention). Role of TFA as ion-pairing reagent is overrated in my opinion, particulary if you are at high organic concentration.
0.1% TFA will give you pH around 1.9 same amount of HCl will give you lower pH (never tried HCl due to corrosion danger). You can try to use ammonium acetate or formate or weaker acid (formic or acetic) and then remove it by distillation or just by stirring with excess of HCl (i did this one successfully when completely replaced mandelic acid for HCl using 4 N HCl in dioxane) in non polar solvent like THF.

regards,

Vlad

Thanks a lot for thoose information !
regards

Nicolas

Surprise Vlad,
full agreement here on "overrated" TFA....
Do you have concrete examples on this "particularily .... at high organic ..."?
Seems to me that I suggested somewhere that I would expect organics + other acids be a possible replacement for TFA.

One more alternative approach: rework the separation to use formic acid. Collect the product. Add a slight excess of HCl. Evaporate the solvent and the product will be in the hydrochloride form. This avoids using HCl in the HPLC system, and you don't have to buy special corrosion resistant equipment. (Of course, if you do want corrosion resistant gear, my employer would be happy to set you up.)

"Neither 0.1% TFA nor 0.01M HCl are a buffer. However they both have sufficient concentration of H+ that they do maintain the pH at around 1. "

By now most of you now know that I have a mission in life to correct the incorrect teaching of buffers common in US courses. So once again(and I thank Mark for the opportunity to teach this again).......

The first of these contradictory statements is wrong and the second is right.

A buffer is something that resists pH changes in response to small additions or acid or base. Mathematically, buffer capacity is defined as the first derivative of the titration curve of the species in question. It follows that where ever a titration curve is flat, buffering capacity is high. There is no combination of 0.01M weak acid and base that will provide as much buffering capacity as 0.01M HCl will deliver at the pH of a 0.01M HCl solution. The classic Clark and Lubs Buffers, for those classically schooled, are preapred from non other than HCl in the pH range of 1 to 2.2.

I hope everyone is convinced now that both strong acids and strong bases can be excellent buffers in their appropriate pH range.

But, really: The buffering at a pH of ~2 (and lower) or ~12 (and higher) is done by the water. (We have been through this before somewhere)
Buffering means that the change of pH upon adding an acid or a base does not change by the increment of activity of added acid or base. This means some H+, etc., has to be "soaked" up by the medium. The already present totally dissociated acid (etc.) can not do this, the water can at the noted pH´s. The lower (higher for basic solution) the pH, the higher the buffering by water. Since 0.01M HCl has a pH of ~2, while 0.01M acetic is higher the buffering of the HCl is higher than that of the acetiic. As I pointed out before: This is especially true if you add acid. If you add base you would quickly get into a pH region where the acetic is a better buffer than the HCl of Bill´s example.

I somewhat agree with Bill, and I taught buffers accordingly. In order to stave off confusion, I added to Bill's definition by stating that for the course purposes buffers were also mixtures of weak electrolytes and their conjugates. I further emphasized that these tended to be the best buffering systems in and around neutrality where living systems tend to proliferate.

Buffering are wonderfully complex yet readily controlled chemical systems. I am glad that chemists have the guts to continue to teach these at the first year college level. Yes, some instructors soft-soap buffers and thereby propagate misconceptions. In my experience, most instructors tried to do it right, as I tried. It was a worthy challenge for all concerned.

Personally, I just don't think of 0.01 HCl as a buffer. I do understand that it resists pH changes under certain conditions. It just doesn't have the equilibria complexity that I think of as a buffer. It is just a strong acid.

FWIW

HW, are you saying that an increment of HCl added to a 0.01M solution of HCl is solvated different from the other HCl in this solution, which results in buffering, or that buffering by HCl is a result of activity coefficient?

All your buffer definitions are too limited and/or too incumbered by mechanism of action. Just as pH is an operational definition with no implied fundamental significance or mechanism, buffer is an operational definition which follows from buffer capacity, d(deltaC)/d pH (where delta C is an increment of strong acid or base). d(deltaC)/d pH is simply the first derivative of titration curve.

To quote from Kolthoff who some claim to be the father of the field of analytical chemsitry, "A solution that resists changes in pH on addition of strong acid or strong base is termed a buffer". Note, this definition implies NO particular mechanism for buffer action. O.01M HCl is a good buffer at its pH simply because a small increment of acid or base doesn't change the total very much, because the total is big relative to the increment. Completely different buffering mechanism from a weak acid buffer, but no less a buffer according to Kolthohff, Clark and Lubbs, and others who developed these concepts.

I used the somewhat sloppy but memorable definition of a buffer:

A buffer is a buffer if it buffers the pH. If it doesn't, it isn't a buffer.

This agrees with Bill's statement that -
"a buffer is something that resists pH changes in response to small additions or acid or base."

Bill, this part of your statement is what I meant:
"HW, are you saying that an increment of HCl added to a 0.01M solution of HCl is solvated different from the other HCl in this solution, "

I am insisting on explaining what goes on chemically since that way one does not forget that adding base to a fully dissociated acid may go to a nonbuffering pH very quickly, while any solution at ~pH = 2 goes toward an extreme buffering capacity on the addition of acid (on decreasing the pH).
Now the dC/dpH = 2.3(Kw/[H+] + [H+] + CKa([H+])/(Ka + [H+]^2))

where Kw is the dissociation constant of water, Ka that of the acid, C is the amount of acid or base added to the buffer. With a large H+ the last term is negligible. So the equation also nicely shows that it´s the Kw that is behind the buffering at low or high pH.

It just flashed into my mind that I might have made a certain mistake in writing the above equation, sure enough. Here is the correct one (from Butler, Ionic Equilibrium):

dCx/dpH = 2.3(Kw/[H+] + [H+] + CKa[H+]/(Ka + [H+])^2)

dCx is the amount of added material, C is the concentration of HA + A- (or ~ the 0.02M of HCl in Bill´s example), also it should be pointed out that the Kw/[H+] +[H+) term derives from the H2O equilibrium.

The excellent point you raise deserves further discussion. It is fresh in my mind as I intend to add this insight to the course I'm preparing for Pittcon.

The buffer capacity (dCx/dpH) can not only be misleading in understanding PRACTICAL strong acid/strong base buffering, but also weak acid buffering as well. (please excuse CAPS, I don't know how to impliment italics) The buffer capacity definition is a mathamatical derivative. In reality, a buffer challenge is an increment of acid or base vastly larger than the imaginary differential amount resuting from the derivative.

Buffer capacity is best understood from examining the titration curve of the species in question. Buffer capacity for additing an increment of strong acid or base is NOT SYMMETRICAL for a strong acid buffer, OR a weak acid buffer, as implied by the differential. That is, the change in pH is, in practice, not the same for adding XX moles of strong base vs strong acid to the buffer, except in the rare case of a weak acid exactly at a pH equal to its pKa.

But the fact remains that Kolthoff has blessed us all to call HCl a buffer at low pH, while we keep in mind how buffers work, least we get burned by a lack of understanding.

I gain insight on this topic every time we argue. I hope we don't hoplessly confuse the audience.

There should be no confusion if one reminds himself of the meaning of dx/dy and if one looks at buffer index curves and the underlying equations.