Erm, this is going to be lengthy and tedious... be warned!
in pH, "p" stands for "potence" (and H is obviously hydrogen ions, or protons). pH is, numerically, the inverse (positive <-> negative, is it "inverse" in english?) of the logarithm of the H+ ions molar concentration. (Or, approximately, the inverse of the exponent of 10 in the concentration).
Example: if H+ ions are 1*10^(-7) Molar, then pH is 7 (which is neutral pH). If we have less ions, then the solution is basic: let's say we have [H+] = 1*10^(-11) M, then pH is 11.
Similarly, pKa is the inverse potence of Ka, which is the dissociation constant for the equation [HA] -><- [A-] + [H+]. So if Ka is 1*10^(-5), then pKa is 5.
An aminoacid has a dissociation constant, so it has a pKa. A solution has a proton concentration, so it has a pH.
Buffer solutions basically work because the buffering agent has (at least) two forms, one of which can take H+ or OH- ions from the solution, thus becoming the other form.
Let's say, acetic acid (AcOH) has two forms, the acid AcOH and the acetate ion AcO- . One relationship between them is that [AcOH] = [AcO-] + [H+], this meaning that if you add protons in the solution (e.g. from a strong acid), then AcO- should take part of them to become AcOH (thus buffering the increase of H+). On the other hand, if you put OH- ions in the solution (e.g. from a strong base), then the OH- will tend to withdraw H+ (becoming H2O together), and AcOH will shift to AcO- making more protons available (thus buffering the decrease of H+). So, the acetic acid form buffers an addition of bases to the solution, while the acetate form buffers additions of acids.
As for this acid, Ka = [AcO-] [H+] / [AcOH],
when [H+] = Ka (that is, when pH = pKa), then [AcO-] = [AcOH].
This means that if pH is equal to pKa, there is an equal quantity of acetic acid and acetate ions.
This is the best condition for buffering activity, because if you add a strong acid, there is a good amount of AcO- to buffer it, and if you add a strong base, there is a good amount of AcOH to buffer it.
But as the two forms are strictly related to [H+], and thus to pH, any increase or decrease of one pH unity will change the concentration of the forms roughly by one potence of 10, or ten-fold!
So, at pH = pKa, a solution of acetic acid / acetate will be 50% acetic acid and 50% acetate... while at pH two units higher than pKa (thus more basic), you will have approximately 1% acetic acid and 99% acetate!
This last solution would be, of course, a risky buffer because it would only buffer acids: to buffer bases, you would need AcOH to be transformed, but AcOH is only 1% of total... so it won't last for long, and it won't work fully either! (As buffers are "weak" acids and bases, there [must] always remain some trace of either form).
pKa, being related to Ka, measures the tendency of an acid to dissociate into ions. So in water, where ions are stable enough, acids have few problems in dissociating. But if you put an acid in a hydrophobic environment, this solution won't host the charged ions as well as water, and acids will tend to retain protons instead of dissociating - thus having a smaller dissociation constant (Ka) and a higher pKa.
*phew* i hope i didn't put mistakes in this!