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Buffers and Titrations

Buffers and Titrations. Dr. Harris Ch 21 Suggested HW : Ch 21: 1, 7, 13, 29. 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

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Buffers and Titrations

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  1. Buffers and Titrations Dr. Harris Ch 21 Suggested HW: Ch 21: 1, 7, 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 NOTneutralize 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 • 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. pH slightly decreases pH slightly increases Initial Buffer solution Base Added (NaOH) Acid Added (HCl) HF HF HF F- F- F-

  5. 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 • O2 is carried throughout the body by binding with the hemoglobin protein found in red blood cells. • Hemoglobin, Hb, binds both H+ and O2. These binding processes are represented below • As expected by LeChatlier’s principle, when blood reaches oxygen-poor tissue, the reaction shifts left, releasing O2 in that region. ΔH > 0

  7. Blood As A Buffer • Under high exertion, the concentration of acid in your body increases for a variety of reasons. • As your heart rate increases, the concentration of O2 in the blood is decreased • High CO2 levels lead to formation of carbonic acid. • Body produces energy by breaking down glucose to pyruvate and ATP. Under anaerobic conditions, lactic acid is formed as a side product from the reduction of pyruvate in O2 anaerobic condition

  8. Blood as a Buffer • The primary buffer systems in human blood are a carbonic acid/bicarbonate buffer system and the HbH+/HbO2 system • A sudden decrease in pH triggers the brain to increase the breathing rate, releasing CO2 faster and shifting the equilibrium right to consume H+. • The body consumes more O2 which shifts the hemoglobin binding equilibrium left, removing more H+. • To operate as an effective buffer, the ratio of HCO3- to H2CO3 must be at least 20:1 conjugate acid/base buffer system

  9. 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 Ka of formic acid is 1.8 x 10-4

  10. 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 base acid neutralized!!

  11. Example • Ex. You have a 100 mL of a buffer that is 0.10 M in CH3COOH and 0.10 M in NaCH3COO 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, CH3COO- • There are exactly .001 moles of HCl added to the buffer, so exactly .001 moles of CH3COO- will be neutralized, forming exactly .001 moles of additionalCH3COOH

  12. Example, contd. Before addition of 10 mL of 0.10 M HCl After .011 mol .01 mol .009 mol .01 mol CH3COO- CH3COO- CH3COOH CH3COOH 100 mL 110 mL .001 molHCl

  13. 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

  14. Group Example • You have 100 mL of a buffer that is 0.3M NH3 and 0.45 M NH4Cl. 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?

  15. 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?

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