1. Buffering Keeping the pH of a Solution Constant (Nearly)

2. Buffer  A combination of a weak acid and a strong base  HAc ↔ Ac - + H +  In the range of pH where the two forms are present H + added from acid will taken up to form HAc H + consumed by addition of base will be replenished from HAc  A combination of a weak base and a strong acid  H 2 CO 3 ↔ HCO 3 - + H +  In the range of pH where the two forms are present H + added from acid will taken up to form H 2 CO 3 H + consumed by addition of base will be replenished from H 2 CO 3

3. Behavior of Vinegar pH of solution barely changes when base added in the range of 3.8-5.8

4. Behavior of Bicarbonate pH Changes little with addition of acid in the range of pH 5.4-7.4

5. There Must Be a Formula! When pH 3-11, this simplifies to Where β max is the maximum amount of acid or base that can be absorbed given some concentration, c, of buffer.

6. Don’t Sweat the Formula  Buffers work best ± 1 pH unit from their pKa  If you expect acids to form, starting pH should be pKa+1  If you expect base to form (acid consumed), starting pH should be pKa-1  A 10 mM buffer can “absorb” about 5 mM acid

7. Volatile Buffers  Sometimes we want to evaporate all the water to concentrate a product like DNA or protein  If the Buffer used is non-volatile like Tris, the salt is left behind  If the buffer components are volatile they will evaporate as well, leaving behind a “clean” product Acetic acid and formic acid are good acids Pyridine and ammonia are examples of good bases to use