Chromatography Forum

I am at a loss here. How does anybody weigh solid KOH or NaOH? How do you titrate accuratly to the desired pH? One can not put an electrode in the solution and even less an indicator.
(I used to prepare the buffers, prior to Bill, with an equally concentrated base and acid version of the buffer, mix them until the desired pH was attained, measured by taking a portion into a test tube. One time it took almost 2 hours to get there, overshooting countless times in either direction. Note that overshooting with this method does not change the desired concentration).

Dear all,
I think I should stop reading all buffer preparation methods now because I totally get lost. I finished with phosphate buffer problem by using KOH and pH meter again because I can't use KH2PO4 for my buffer ( I must use 80% MeOH for my mobile phase). However, another buffer is making me headache - 0.1M acetate buffer pH 4
In "HPLC method development", they prepare acetate buffer pH 4 by mixing 820ml of 0.1M CH3COOH and 180ml of 0.1M CH3COONa. The recipe from http://www.bi.umist.ac.uk/users/mjfrbn/ ... akebuf.asp calls for 0.084mol of CH3COOH and 0.015 mol CH3COONa. And from http://www.zirchrom.com/Buffer.asp, they mix 1000ml of 0.1M CH3COOH and 174ml of 0.1M CH3COONa to prepare 1000ml of 0.1M acttae buffer pH 4
I tried the first way, the buffer showed pH 3.2
The second way, pH was 3.6
The third way, I used higher concentration of CH3COOH and CH3COONa, both were 1M, and I mixed 100ml of 0.1M Ch3COOH and 17.4ml of 0.1M of CH3COONa (recipe form zirchrom website), I got the pH 3.7.
I also tried to mix them with ACN for my mobile phase and at last I chose the way from Zirchrom. However, I am still not sure how to report the buffer prepration method. Should I report 0.1M acetate buffer pH 4 or pH 3.7?
Sorry to bring alot of trouble to all of you but this kind of discussion is really useful for a student like me!

Do the whole thing again, but this time make sure that your ionic strength and temperature are the same for the recepies, that your reagents are OK, and that the pH meter is calibrated correctly (~no drift, etc.).
Didn´t want to bother with the login at zirchrom, but 1000mL + 174mL of those reagents don´t make 1L.

Oh yes, too bad you didn´t try a lower concentration of phophate buffer (prepared by weighing the acid, base forms...).

When using potassium hydroxide or sodium hydroxide as one of components used for preparing a buffer, we use the appropriate standard concentrated solution (i.e. for sodium hydroxide this is a nominal 50% solution in water). It's not necessary to titrate such a solution as quality analytical reagents of this sort have a label assay accurate to better than one part in 500. One simply weighs out the appropriate amount of the concentrated solution based on label assay with the required stoichiometry.

I guess it's possible that using prepared salts might be comparable in accuracy to weight based in preparation of buffers from acids and bases but assays on acids and bases are typically accurate to 1 part in 500 to 1 part in 1000. On the other hand, the salts are not necessarily prepared with such accuracy. In our experience, preparing by weight using only acids and bases is more precise.

And these "standard" solutions are useable for UV, fluorescence detection? How do the manufacturers calibrate them? (Is Bill right on them using solids.....?).
Also, how long are these hydroxide solution "standard", especially after you open them? (I have a bottle in the fridg, over one year, didn´t need it as it was planned to use it only if no other way was viable, should probably throw it away now?)

50% sodium hydroxide concentrate is quite stable although when using it for reagent purposes it's advisable not to buy too large a container. Several suppliers sell it in 25 ml containers which cost in the range of $15. This size is convenient for making buffers. Fluka sells a grade specifically intended for HPLC applications although this grade is not available in containers less than 1 l. Anyway, I wouldn't keep it around for long period of time if you're planning on using it for making quality buffers anymore than I would want to use an old bottle of phosphoric acid for use in preparing buffers. It's probably not a great idea to store concentrated sodium hydroxide in the refrigerator, though. There is nothing about the reagent that will induce degradation at room temperature, nothing can grow in 50% sodium hydroxide but cold sodium hydroxide will pick up moisture more readily than hydroxide stored at room temperature.

Cold NaOH picks up moistutre if you open it while still cold.
I put it in the cold room to reduce dissolution of monomers, plastisizers... from the plastic, maybe even CO2 diffusion (I know HCl goes through polyethylene from discolored labels on standard HCl solutions) through the plastic container. I have, yet, never used solvents, liquids, from a plastic container for HPLC. I wouldn´t even consider concentrated hydroxide solutions from class containers.

I investigated the Zirchrom approach to buffer preparation in 2002, so what follows is true as of that date. Turned out I knew the person that wrote the software. Their buffer calculator did not account for activity coefficients, and hence, it is useless as a guide to preparing a buffer to some target pH. If you ignor activity coefficients you are bound to get the wrong pH from their recipe.

I don't understand all the fret over pH. Make up something close, by weight and volumn of course so that the preparation can be repeated precisely, and try it. If the separation works, next time repeat buffer preparation as before and it will work again. Report weights and volumns and others can precisely repeat the preparation. The separation is a much better indication of "correct" than the pH meter. I worked in an LC lab for 8 years before I bothered owning a pH meter and then I bought one for another project.

Specifically, make up the acetate buffer by whatever recipe you prefer. Test the separation. If it works who cares what a meter says it to be.

In response to HW: Sodium carbonate is not soluble in concentrated sodium hydroxide. Postasium carbonate is soluble in its concentrate solution. These facts have different ramifications depending on the nature of eluant/buffer one wishes to prepare. The choice is obvious for "carbonate free" ion chro. eluants. CO2 most definitly will diffuse through plastic bottles. That is why one does not see small volumn carbonated beverage bottles(surface to volumn ratio). Poly(ethylene) does not have plasticizer. It may have UV and antioxidant stabilizers at very low concentrations, but they are not readily extracted by water.

Like Chris I always kept a few ampules of "standard" 1M sodium hydroxide and 1M HCl concentrate around for preparing certain solutions, for example Clark and Lubbs buffers. I never experience any LC problems using these standardized ampules of concentrate. But when I could, buffers were preferably prepared from salts.

Ups, glass not class was meant above.
Bill, assuming that you are talking of UV detection it is good to know that one probably has a good chance not to get a mess of peaks due to the containers.
I used to make acetate buffers by getting a HPLC grade acetic, preparing the desired conc aqu solution, splitting it, quickly popping some NaOH pellets into one part and mixing the two solutions until the desired pH was attained.
Looks like standard NaOH solution is worth a try.

Concentrated sodium hydroxide will absorb carbon dioxide. Bill is right about sodium carbonate not being soluble in concentrated sodium hydroxide, the carbonate is the white precipitate you see in a bottle of soduim hydroxide exposed to air. Remember the bad old days of Ascarite as a gas purifier? it was used to remove traces of water and carbon dioxide. Composition - sodium hydroxide on asbestos.
The main objection I have to the use of sodium hydroxide is the variable, and usually relatively low. purity. which changes with time due to formation of sodium carbonate on the surface of the pellets. Using solutions from a sealed ampule will help in this regard, but I still prefer using salts.
Bottom line here is that there is no one correct way to prepare a buffered solution. I may use one method, you may use another, and we both can get good, reproducible results. The problem comes in if I try to make a buffered solution to a certain pH you have specified in a method, and I make my buffer using a different technique than you do. In that case we may not get equivalent results. That's the reason I like to see the method used to prepare buffers specified in a method.

While I too appreciate the advantages of Bill's buffer prep. methods, I must interject one more comment. If a method delivers poor resolution as a result of a trivial difference in pH, it is not a very robust method. In such cases, buffer prep. should be worked out so that, in the worst cases, weights of dried reagents can be specified in the method in a detailed manner.

I also appreciate that 0.3 is a pretty large difference in pH and I would also suggest that there is likely some procedural/training issue that needs attention here.

Thanks, Ron and Dr for summarizing, but, Ron this NaOH pellet stuff had a bit other thinking behind it. First, to make things clear at the outset, I am convinced enough on standard hydroxide solutions to try that, BUT only if adding hydroxide is the only way to get at the goal. Before Bill I prepared phosphate buffers, as an example, only via weighing base and acid form at the desired concentration into two separate 1L volumetric flasks, etc., then mixing the "acid" and "base" solutions until the desired pH was attained (sometimes tedious, but I NEVER used hydroxide here).
Now back to the earlier use of NaOH pellets. The thought of just guessing ("educated") the amount of pellets instead of weighing them was to reduce the H2O-CO2 uptake drastically, and rendering what was picked up inconsequential. The pellets were added to aqu. acetic, so any adsorbed (on the pellets) H2O didn´t bother anything. It was also figured that the CO2 would come to the same equilibrium in the final solutions, whether some was adsorbed on the pellets or not. (Remember, I didn´t weigh the pellets!!). Now, in the absence of standard hydroxide solution I still think this is a viable, reasonably reproducible technique.

HW, I agree with the logic you used with the hydroxide pellets. The pellets are generally uniform in weight, and the less time exposed to air the better. In the absence of a standard preparation the less handling of the pellets the better. You are probably getting better reproducibility in the preparation than if you spent a significant amount of time trying to weigh out the same mass time after time.