Worksheet 21 - Buffers
Until now, we’ve been primarily concerned with calculating the pH of a solution in which
we initially added only one acid or base. Now we will consider solutions that initially
contain both a weak acid (HA) and its conjugate base (A-).
Suppose a solution contains both HF and NaF. What are the major species in solution?
We know that HF is a weak acid (Ka = 7.2x10-4), so the major species resulting from HF
are HF and H2O. NaF is an ionic compound, so it dissociates in water to form Na+ and
F- (and of course H2O is still a major species in this solution). Altogether, the major
species are HF, Na+, F- and H2O.
HF + H2O ' F- + H3O+
The weak acid produces some F- and H3O+ in solution. By adding NaF, we’re
introducing another source of F-. LeChatelier’s principle indicates that the equilibrium
should shift to the left upon addition of a product. This suggests that the [H3O+] will
decrease, and the pH will increase. A solution containing both a weak acid and its
conjugate base should be less acidic than a solution containing only a weak acid.
A buffer solution is one which maintains an approximately constant pH when small
amounts of either a strong acid or base are added. Solutions containing either weak
acids and their conjugate bases, or weak bases and their conjugate acids, can be
buffering solutions.
Here are some important facts about buffer solutions:
• The number of moles of the conjugate pairs must be large compared to the
moles of added strong acids or bases
• The ratio of the conjugate pair ([weak acid]/[conjugate base] or [weak base] /
[conjugate acid]) should lie between 0.1 and 10, with optimal buffering at a 1:1
ratio
• The pH of a 1:1 ratio buffer is equal to the pKa of the weak acid or pKb of the
weak base. The effective range is ± 1 from the pKa or pKb.
The Henderson-Hasselbalch equation is particularly useful for calculating the pH
of buffer solutions:
⎟
⎟
⎠
⎞
⎜
⎜
⎝
⎛
+= ]acid[
]base[
log
a
pKpH
(
)
()
bb
aa
KpK
KpK
log
log
−=
−
=
