1 Rev 1/2013 Experiment #5 -pH Buffers Required Skills To perform this experiment you need to have mastered the entire laboratory skills required for Experiment #3. Most importantly, you need to have developed a solid theoretical understanding of pH buffers. A short summary follows, but each student should read and understand the sections in their CHM 158 textbook regarding buffers. If the background material is mastered prior coming to lab, this lab is very easy. Background A buffer is a solution that has an appreciable concentration of a weak acid and its conjugate base. Such solutions are called “buffers” because they resist change in pH when the volume of the solution is changed (through either dilution or concentration) and when small amounts of strong acid or strong base are added. Most naturally occurring aqueous solutions (e.g., blood and lakes) are buffered. In the laboratory, buffers can be created in numerous ways, but the three most common ways are: i. Adding comparable amounts of the weak acid and its conjugate base to water. ii. Adding a limited amount of strong acid to a solution of a weak base. iii. Adding a limited amount of strong base to a solution of a weak acid. Predicting the pH of a Buffer The pH of a buffer solution can be easily derived from the equilibrium expression associated with the dissociation of the weak acid: HA + H2O ? A-+ H3O+ (Rxn 1) [ ][ ] [HA] 3 + a K = A H O – (Eq 1) Equation 1 can be rearranged to yield [ ] [ ] [A- ] H O = Ka + HA 3 (Eq 2)
CHM 158: Experiment #5 2 If we take the negative logarithm of each side of equation 2 we obtain [ ] [HA] log A- pH = pK + a (Eq 3) where pH = -log[H3O+] and pKa = -log Ka. Eq. 3 is known as the Henderson-Hasselbalch equation, and it is the fundamental relationship used to predict buffer pH from the relative concentrations of the acid and its conjugate base. The Henderson-Hasselbalch equation can also be expressed in terms of the moles of the weak acid and the conjugate base ( ) (moles HA)…