Question

Another question...

I just had one more thing to add- I'm having trouble understanding, conceptually, why a strong acid-strong base combination makes a poor buffer system. When you compare titration curves of a strong HA/strong A- with strong HA/weak A- (or weak HA/strong A-), they look the similar to me...they both have areas before (and after, actually) the equivalence point where pH does not change rapidly, but only in the weak/strong combination is this considered a buffer range. Why is this? I think the answer has something to do with the purpose of a buffer (to neutralize added acid or base), but I'm having a hard time understanding how a strong acid-strong base combination fails to do this... Also, I found this flash web-page, and personally I find it helpful to visualize these sorts of things, thought you might want to pass it on to the rest of the class: http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/buffer12.swf

This really gets to the root of why and how a buffer works. Buffers control pH because there is a significant quantity of both the conjugate acid and conjugate base present in solution to react with added acid or base. Let's put some numbers on it to convince ourselves. Let's say we have a solution that contains 10,000 molecules of HA and 10,000 A^- ions. If I throw 50 H⁺ ions into this solution, the ratio will be 10,050:9,050. It has changed, but it hasn't changed much so the pH of the solution remains nearly constant.

Now let's think about a strong acid solution. In that strong acid solution, the concentration of conjugate acid is effectively zero because strong acids (essentially) completely dissociate, so the ratio of HA to A^- is more like 1:1,000,000 (or more). If you add a little OH^- to this mixture, it's not going to react with HA because there's so little of it present. It will react with H₃O⁺ that's floating around free in the water and the resulting solution will be influenced not by the HA/A^- equilibrium, but by the H₃O⁺/H₂O/OH^- equilibria present in water. This is the pH leveling effect of aqueous solution. It is possible to make solutions that have negative "pH" or "pH" above 14, but these are not regular aqueous solutions and the definition of "pH" has to be stretched a little bit to understand the acid/base character of these solutions. Convinced? Pull out your calculator and calculate the pH of a 6M aqueous solution of a strong acid. If we assume that 6M strong HA results in 6M H⁺ ions (or H₃O⁺ ions if you prefer), the pH of that solution would be -0.78. What about a 1M solution? It would have a calculated pH of zero. In practice, these are not the measured pH's of these solutions because the ionization/autoionization of water kicks in and limits the observed pH.

Thanks for the web site, it's got a few other flash animations that might be helpful.
http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/flash.mhtml