Zeb Karimi

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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(()/())#

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 ## in order to determine pH, since

#pH = -log()#

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 ## = ##, which means

#pH = -log()#

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 = ( * )/()#, which means that you'll get

#K_a = (x * x)/(C-x) = x^(2)/(C-x)#

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