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So I was wading through several textbooks yesterday reviewing the concepts of buffers. I found them lot more challenging than I used to. One thing in particular disturbed me: The books and internet all said "Yeah, buffers work pretty great in as long as the $\rm pH \approx pK_a\pm 1$." But I am not one to leave these things alone. I finally found the Quantitative explanation I was looking for over at buffer capacity on Wikipedia.

However, the graph inside the link says the buffering capacity becomes great when you solution becomes very acidic or very basic, as well. And, of course, looking at titration curves there seem to be a few "buffer" zones e.g. where the solution is highly acidic or highly basic, the $\rm pH$ doesn't change as much after a certain point:

enter image description here

So, obviously there is something wrong with my reasoning, since nobody calls those areas buffer zones. Is it because it is not useful to have such a highly acid/basic solution? Or is it just because technically, there is no buffering, just small $\rm pH$ changes? Or maybe something else entirely...

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  • $\begingroup$ I totally agree with trb456. I'd like to add that what you're seeing in the figure you provided is actually a titration of water (that starts out at acidic pH). Remember that water is both an acid and a base. When you add NaOH as in your figure, that titration curve is what you see. That's why the inflection point occurs at pH = 7.0. You add enough NaOH to neutralize the nitric acid and get to pH 7, then as you add more NaOH you increase the concentration of hydroxide ion - increasing the pH. Also, remember that pH is calculated on a log scale, which helps explain the shape of these plots. $\endgroup$ – Phillip Mar 4 '13 at 23:33
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I think the issue with buffer systems is that they are resistant to changes from strong acids/bases in their normal state. For example, biological systems often are buffered so that they have some resistance to changes in pH, which can wreck havoc on biochemistry.

Recall the basic buffer equation:

$\ce{HA <-> H+ + A-}$

where $\ce{HA}$ is a weak acid, and there is an excess of $\ce{A-}$, the conjugate base. Because $\ce{HA}$ is weak, it is a relative poor proton donor, but $\ce{A-}$ is a relatively good proton acceptor. So if you now add a strong acid, the $\ce{H+}$ it donates reacts with the $\ce{A-}$ to produce $\ce{HA}$, shifting the equilibrium to the left, rather than to the right as one might expect. Or perhaps, preventing as much shifting to the right as might be expected. This gets to your issue of buffering capacity.

This line of reasoning does not work with strong acids/bases because they donate/accept protons with gusto, and thus are not resistant to shifting the equilibrium. You would need to be way at the extremes of pH for there to be a small change to pH, as your graph shows. In contrast, buffers are resistant to pH change in their normal state.

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Buffering capacity is just the amount of acid (strong enough to protonate everything) you must add per mole to bring about a unit pH shift. For example, the buffering of pure sulfuric acid at its second pK is, theoretically, 0.57 Eq/mol-pH and falls off very rapidly at pH values removed by more that 1 pH - just like any other acid. On the other hand a solution of water with sulfuric acid at mole fraction 0.1 has buffering of 0.08 but does not exhibit a peak at pH 1.99. This is because 0.1 mole of acid has buffering capacity of 0.057 while 0.9 mol of water has buffering capacity of 0.02. The buffering of the water is approaching the buffering capacity of the acid. At higher pH where we make buffers with weaker acids the buffering of the water itself is much smaller and so the acid dominates ad the peaks near the pKs become evident.

This is largely a problem with semantics. We are taught that a buffer is something that holds pH near a particular value by absorbing the 'stress' of added acid or base. When we speak of the high buffering capacity listeners tend to assume that that there is a peak in any case where in many cases there isn't.

At low pH added acid goes to protonate the solvent, not so much the anion of the solvent. The nice Gaussian shaped buffering vs pH curve goes away.

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