I'll discuss how to determine pH given "pKa" for a monoprotic acid, which is an acid that only donates one proton per molecule when placed in aqueous solution.
The general equation for a monoprotic acid in aqueous solution is
HA_((aq)) rightleftharpoons H_(aq)^(+) + A_(aq)^(-)
If you're dealing with a buffer, then you are dealing with a weak acid. In this case, the Henderson-Hasselbalch equation can take you directly from "pKa" to the solution's pH (assuming you know the concentrations of the weak acid and its conjugate base)
pH = pKa + log(([A^(-)])/([HA]))
If you're not dealing with a buffer, then you must use the acid dissociation constant, "K"_a, to help you determine the pH of the solution. In this case, you need to determine [H^(+)] in order to determine pH, since
pH = -log([H^(+)])
The value of the acid dissociation constant can be derived from "pKa"
K_a = 10^("-pKa")
For a strong acid, "pKa" <1 and "K"_a>1 ; strong acids dissociate completely in aqueous solution, so [H^(+)] = [HA], which means
pH = -log([HA])
If you're dealing with a weak acid, you have to use the ICE table method (more here: http://en.wikipedia.org/wiki/RICE_chart). The initial concentration of the acid is C, so
......HA rightleftharpoons H^(+) + A^(-)
I:......C.........0.........0
C:...(-x).........(+x)......(+x)
E:..(C-x).......(x).......(x)
Remember that "K"_a is defined as
K_a = ([H^(+)] * [A^(-)])/([HA]), which means that you'll get
K_a = (x * x)/(C-x) = x^(2)/(C-x)
At this point, the only unknown you'll usually have is x; solve for x and pick the positive solution (remember that x represents the concentration of H^(+) and A^(-), so it must be a positive number).
As a result, [H^(+)] = x => pH = -log([H^(+)])
