(For mobile users) 00:08 Dissociation of (generic) weak acid HA 00:49 Acid dissociation constant Ka 02:55 In terms of [H₃O⁺] (=[H⁺]) 03:30 Taking log of each side 04:02 First property of the logarithm (log ab = log a + log b) 04:31 Applying property 1 05:09 Multiply each side by -1 05:37 Use definitions of pH and pKa 06:10 Second property of the logarithm (log a/b = - log b/a) 06:30 Applying property 2
THANK YOU SO MUCH. My biochem textbook just says 'take the -log and make some rearrangements' and I literally could not figure out how this happened. You have restored my sanity.
Recall that the job of a buffer is to minimize the change in pH when we add either strong acid or strong base to the solution. Suppose we have an acetate buffer. Here, the weak acid (HA) is acetic acid (CH3COOH), and its conjugate base (A-) is acetate (CH3COO-), perhaps from sodium acetate. If we add strong base (say sodium hydroxide), it gets neutralized by HA: NaOH + CH₃COOH ⇒ Na⁺CH₃COO- + H2O. If we add strong acid (say hydrochloric acid), it gets neutralized by A-: HCl + CH₃COO- ⇒ CH₃COOH + Cl- If nothing is added to the solution, the pH depends on the "competition" between HA and A-. The stronger the weak acid HA and the more of it that is present, the lower the pH will be (since a stronger weak acid will have a LOWER pKa, and the smaller will be the log fraction in the H-H). The stronger the conjugate base A- (and the more of it that is present), the higher the pH will be. Does that help?
