Actually, strong acids and strong bases do have a buffer region while the acid/base is relatively concentrated.
Notice on the SA/SB titration that pH changes little with the addition of the titrant (at the top and beginning of the titration) so it acts as a buffered solution. This is not emphasized in most classes but you can see it is so. I think most chemists do not consider this a buffer; however, it follows the definition that a relatively large amount of titrant can be added with only minimal changes in the pH of the resulting solution. One of the profs I had in graduate school pointed this out to me one day while I was making a graph similar to the one in the above link. The effect diminishes as the concn of the acid/base decreases.
so in strong acid strong base titration does the pH=pkA?
it doesn't right because for like HCl it doesn't really have a Ka?
Strong acids and strong bases don't have a Ka or pKa because they are 100% ionized. Weak acids and weak bases have only one point on the titration curve where pH = pKa an that is where the (base) formed = (acid) remaining. When that is true, then
pH = pKa + log(b/a), and when b = a, then the fraction is 1 and log 1 = zero; therefore pH = pKa. Few people will count strong acids or strong bases as buffering agent (that includes me); however, there IS A REGION where they show buffering action as defined by the definition of a buffer.
So then the common answer for any titration with a strong acid there is no buffering/buffer region?
I don't think I said that. I said that few people will count strong acids or strong bases as buffers BUT they do show regions where they show buffering action, especially if they are concentrated (for example in 1 or 2 M HCl) titrated with a strong base. For example, if we take 100 mL of 1M HCl and titrate with 1M NaOH. here is a table. HCl heading = mmoles
NaOH heading = mmoles.
Start with 100 mL of 1M HCl = 100 mmoles HCl.
HCl....mL 1MNaOH....HCl left....pH
......101.00 mL......1 NaOH...11.7
......150.00 mL......50 NaOH..13.3
......151.00 mL......51 NaOH..13.31
I hope this shows you what I mean. You see that at the concentrated levels of both acid and base, the pH changes little with added titrant but near the equivalence point SMALL amounts of titrant make huge changes in the pH. Of course, this is why titrations work so well because of that huge change at the equivalence point.
I understand what you're saying, but in my text book it says a buffer is a weak acid or base with its salt, so is that not always true?
Yes, that is the correct definition for a buffer and it's the standard one. Everyone uses that and you should not stray from it. And it is ALWAYS true. So a weak acid and its salt will be a buffer. A weak base and its salt will be a buffer. Having said that, it doesn't mean that in some circumstances other mixtures CAN'T act as buffers although they don't meet the official definition of a WA/WB and the salt. For example, a mixture of acetic acid and sodium acetate certainly is a buffer. But what about sodium acetate and HCl? With an excess of sodium acetate, the HCl ties up acetate ion to form acetic acid so the mixture of sodium acetate and HCl can act as a buffer. It may not be as effective as sodium acetate and acetic acid but it functions along those lines. I'm not trying to confuse you; I'm just trying to show you the difference between an official buffer and cases in which buffering regions can occur without those "official" ingredients. If that is confusing to you, then just stick with your definition of a weak acid or weak base and its salt. Those will always keep you on the right path.