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The concept of buffers, focusing on solutions containing both a weak acid and its conjugate base. It explains the major species in such solutions, the impact of lechatelier's principle, and the importance of buffer solutions. The document also provides several examples to calculate ph using the henderson-hasselbalch equation and discusses various weak acids and bases. Students will learn how to determine appropriate buffer components and understand the role of weak acids and bases in maintaining a constant ph.
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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 H 2 O. NaF is an ionic compound, so it dissociates in water to form Na+^ and F-^ (and of course H 2 O is still a major species in this solution). Altogether, the major species are HF, Na+^ , F-^ and H 2 O.
HF + H 2 O ' F-^ + H 3 O+
The weak acid produces some F -^ and H 3 O +^ 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 [H 3 O +^ ] 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 Henderson-Hasselbalch equation is particularly useful for calculating the pH of buffer solutions:
⎟⎟ ⎠
[acid]
[base] pH pKa log
a a pK K
pK K log
log = −
Here are some Ka and Kb values for various weak acids and bases (shown with their conjugates): CH 3 COOH/CH 3 COO -^ K a = 1.8 x 10- NH 3 /NH 4 +^ K b = 1.8 x 10- HCO 3 - /CO 3 2-^ K a = 5.6 x 10- H 2 PO 4 - /HPO 4 2-^ K a = 6.2 x 10-