1. BUFFERS AND TITRATIONS Dr. Harris Ch 21 Suggested HW: Ch 21: 1, 13, 29

2. Buffers Solutions that contain both a weak acid and its conjugate base are called buffers. Buffers have the unique ability to resist sharp changes in pH A buffer is able to do this because it contains both acid and base components However, the acidic and basic components of the buffer themselves do not neutralize each other, a special property fulfilled by mixing conjugate pairs If an external acid or base is added to the system, it is instantly neutralized by one of the conjugate species

3. Creating A Buffer For example, an HF/F- buffer can be created by mixing 0.1M HF with 0.1 M NaF. Since NaF will fully dissociate, we ignore the neutral Na +. In solution, you’d have: Although HF is an acid, and F - is a base, the two CAN NOT neutralize each other because the reaction has no preferred direction. Thus, by mixing an acid with its own conjugate base, you create a buffered solution in which the components do not affect each other.

4. How Buffers Work HFF-F- F-F- F-F- Initial Buffer solution Base Added (NaOH) pH slightly increases pH slightly decreases The acidic component of the buffer neutralizes any base that is added to the system. The basic component of the buffer neutralizes any added acid. The resulting changes in pH are very small. Acid Added (HCl)

5. C.I.R.L.: Blood as a Buffer Blood is the best example of a buffer in everyday life. The pH of the human body is 7.4. If the pH of the body becomes to acidic or basic, certain proteins and enzymes are chemically altered, rendering them inactive If the body’s pH drops below 6.5 or above 8.0, death may occur. Human blood acts as the body’s buffer to prevent pH changes

6. Blood as a Buffer O 2 is carried throughout the body by binding with the hemoglobin protein found in red blood cells. Hemoglobin, Hb, binds both H + and O 2. These binding processes are represented below As expected by LeChatlier’s principle, when blood reaches oxygen-poor tissue, the reaction shifts left, releasing O 2 in that region.

7. O O

8. Blood As A Buffer (ex. exercising) anaerobic condition in O 2 Exercise Need energy fast. Must burn more carbs. Breathe faster. Bring in more O 2 Heart rate increases to transport extra O 2 Combustion of carbs yields CO 2 and H+ in the muscles Not enough O 2 ? Anaerobic breakdown of glucose yields lactic acid

9. Muscle Blood H + CO 2 O2O2 Concentration Gradients Shift Equilibria

10. Blood as a Buffer A sudden decrease in pH triggers the brain to increase the breathing rate, releasing CO 2 faster and shifting the equilibrium right to consume H +. The body reacts O 2 faster, which shifts the hemoglobin binding equilibrium left, removing more H +. Lactic acid is neutralized by bicarbonate, HCO 3 -. Blood pH remains intact. The primary buffer systems in human blood are a carbonic acid/bicarbonate buffer system and the HbH + /HbO 2 system

11. Calculating the pH of a Buffer Using the Henderson- Hasselbalch Equation The Henderson-Hasselbalch equation is used to determine the pH of a buffer system. Example: Calculate the pH of a formic acid buffer (HCOOH) that is.25M HCOOH and.15 M of the conjugate base (HCOO - ), given that the K a of formic acid is 1.8 x 10 -4

12. Calculating the pH of a Buffer Solution After Exposure to an Acid/Base Let’s calculate the pH of a buffer solution following a response to a pH disturbance. It is important to know that a molecule of acid completely neutralizes a molecule of base, forming the conjugate base and conjugate acids So, to determine changes in concentration of either the acidic or basic component of a buffer, simply subtract the number of moles that are neutralized from the active component, and add those to the inactive one acid base neutralized!!

13. Example Ex. You have a 100 mL of a buffer that is 0.10 M in CH 3 COOH and 0.10 M in NaCH 3 COO at a pH of 4.74. If 10 mL of 0.10 M HCl is added to the buffer solution, what is the new pH? What is the change in pH? Since HCl is an acid, it will only react with the basic component of the buffer, CH 3 COO - There are exactly.001 moles of HCl added to the buffer, so exactly.001 moles of CH 3 COO - will be consumed, forming exactly.001 moles of additional CH 3 COOH

14. Example, contd..001 mol HCl Before addition of 10 mL of 0.10 M HCl CH 3 COOH CH 3 COO -.01 mol 100 mL CH 3 COOH CH 3 COO -.011 mol.009 mol 110 mL After

15. Example, contd. The new concentrations are: Buffer works well. This is a very small drop in pH considering that HCl is a strong acid. ΔpH = -.09 * Since both the acid and base are in the same volume of solution, you may substitute moles for concentrations in the equation. The ratio is the same.

16. Group Example You have 100 mL of a buffer that is 0.3M NH 3 and 0.45 M NH 4 Cl. The pH of the buffer is 9.07. Find the pKa. Calculate the change in pH after the addition of 5 mL of 4.0 M NaOH?

17. Group Example You have 150 mL of a 1M HCl solution. You add 150 mL of a 1.2 M NaOH solution. What will the pH be? (This is NOT a buffered solution)