Hi all,
Just had a quick question about for most of the buffer system, why is making buffer at pH pKa±1 is recommended? What is the reason behind it? I've seen people violating this rule all the time.
For instance, people use 200 ammonium acetate at pH 7.0 for protein analysis using mass spec.
I've seen a ton of other cases.
why is making buffer at pH pKa±1 is recommended?
Dear Cece Hong
Hi,
For calculation of pH for a buffer solution we should use equation pH = pKa + log[b]/[a]. When [b]/[a] is 10, then one unit will be added to pKa in the above equation. In addition, when [b]/[a] is 1/10, we will have pH = pKa - 1. Based on this equation, we concluded that any buffer has the most power to show constant pH.
The answer is in textbooks.
" I've seen people violating this rule all the time. " This indicates that these people have no basic knowledge in chemistry. I've seen numerous papers in prestigious journals violating this simple rule. It's waste of time reading such papers
from https://www.merriam-webster.com/dictionary/rule
"To play this game (chemistry in this case), you need to follow the rules."
The ability of a buffer solution to maintain a nearly constant pH when a small amount of acid or base is added to the solution is greatest at the pKa and diminishes as the pH of the solution goes above or below the pKa. A rule-of-thumb is to use a buffer within 1 pH unit of the pKa to maximize its buffering capacity (see the Henderson-Hasselbach equation provided by Aziz Habibi-Yangjeh for the math).
However, if there is a desire to use a particular buffer outside of its preferred pH range, the buffering capacity of the solution can be increased by increasing the concentration of the buffer. Ammonium acetate (pKa 4.8) (also ammonium bicarbonate and ammonium formate) is used for mass spectrometry because it is volatile. Such a high concentration as 200 mM ensures that the pH will not change unless a substantial amount of acid or base is added. If no acid or base is going to be added, then the pH will not change from where it started, so the weak buffering isn't a problem.
200 mM ammonium acetate buffer at pH 7.0 has a very low buffer capacity. An addition of 0.5 mM acid or base will shift pH by about 0.1 unit. 200 mM Sodium acetate buffer has lower buffer capacity than ammonium acetate. An increase of buffer concentration significantly changes the ionic strength, which might affect equilibria more than effect of pH. In general, as Adam correctly stated that "If no acid or base is going to be added, then the pH will not change from where it started, so the weak buffering isn't a problem."
The effective buffer pH range for weak acid / strong base can be understood as follows:
I. For a not too dilute weak acid solution (approx. Ca > 2.5·103·Ka) we can predict that (*):
pH = ½(pKa + pCa); with pCa ≡ - log10(Ca/M)
II. Let that acid be titrated with a strong base. At the titration end point, only salt is found in the titrated solution, and we can predict that (*):
pH = ½(pKw + pKa - pCs); with pCs ≡ - log10(Cs/M)
III. The pH range covered along that titration, from pure acid up to the titration ending point, would be: ΔpH = ½(pKw + pKa - pCs) - ½(pKa + pCa) = ½(pKw - pCs - pCa)
IV. Within that range, the buffering effect is most effective close to the pH where the acid is half-titrated. The pH varies quite steeply when the titration ending point is approached; and not so steeply, but still significantly, close to titration start. It is quite reasonable to conservatively take, as effective buffer range, an intermediary pH range of 2 units from the pH range covered along the titration.
V. The Henderson–Hasselbalch equation predicts the pH within the above mentioned buffer range (*): pH = pKa + log10(Cs/Ca). When the acid is half-titrated (Cs = Ca): pH = pKa. The effective buffer range is thus given by: pH = pKa ± 1
(*)https://www.researchgate.net/post/pH_calculation_of_a_mixture_of_formic_acid_NaOH_and_water
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