1 / 15

Buffers

Clinical Analytical Chemistry CLS 231. Buffers. Lecture 3 Lecturer: Amal Abu Mostafa. Session Objectives:. Buffer solution Composition of a Buffer Examples of buffers How a Buffer Works Preparing Buffers with a Desired pH Henderson Hasselbalch equation

lucian-richardson
Télécharger la présentation

Buffers

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Clinical Analytical Chemistry CLS 231 Buffers Lecture 3 Lecturer: Amal Abu Mostafa

  2. Session Objectives: • Buffer solution • Composition of a Buffer • Examples of buffers • How a Buffer Works • Preparing Buffers with a Desired pH • Henderson Hasselbalchequation • Solving Problems Using the Henderson- Hasselbalch Equation

  3. Buffer solution: • Buffer Solution: Solution that resist change in the pH because it contain both acidic and basic species to neutralize OH- and H+ respectively. • Buffers are aqueous systems that tend to resist changes in pH when small amounts of acid (H+) or base (OH-) are added. • Composition of a Buffer: • Buffer consist of a weak acid and a salt containing its conjugate base. • Buffers may also contain a weak base and a salt containing its conjugate acid. • However, that acid and base must not be able to neutralize each other.

  4. Examples of buffers: • - acetic acid / sodium acetate • CH3COOH/ CH3COO– • - NH3 / NH4+ • - Blood contains buffers that maintain a consistent pH of about 7.4 , so that biochemical reaction can follow their correct path. • Blood contain carbonic acid and bicarbonate ion (its conjugate base) H2CO3/HCO3- buffer that help to maintain pH=7.4

  5. How a Buffer Works: • To work, a buffer must be able to neutralize either a strong acid or strong base that is added. This is precisely what the conjugate acid and base components of the buffer do. • For example: CH3COOH/ CH3COO– • If we add extra H+ to the buffer (from a strong acid) the acetate ion (the weak conjugate base) can react with it as follows. • H+ (aq) + CH3COO– (aq) → CH3COOH (aq) • A similar response occurs when a strong base is added to the buffer. The OH– from the strong base will react with some CH3COOH • CH3COOH (aq) + OH – → CH3COO– (aq) + H2O

  6. The equilibrium position of the buffer is governed by the reaction • CH3COOH(aq) + H2O(l) →H3O+(aq) + CH3COO– (aq) • In a general form: • HA H+ + A- • And its acid dissociation constant • Ka = [H+][A- ] = [H+](mol A- . L- 1 ) = [H+](mol A-) • [HA] (mol HA . L- 1) (molHA) • This means that for a given acid-base pair, [H+]is determined by the mole ratio of conjugate base to conjugate acid, we don’t have to use molar concentration. • For buffer solutions only, we can use molar concentration or moles in the Ka (or Kb) expression to express the amounts of the acid and its conjugate base.

  7. Preparing Buffers with a Desired pH • The hydrogen ion concentration of a buffer is controlled by both Ka and the ratio of concentrations (or ratio of moles) of the members of the acid-base pair. • Ka = [H+][A-] • [HA] • By rearranging this equation to solve for [H+] • [H+] = Ka [HA] (1) • [A-] • Solutions of conjugate acid-base pairs will generally act as buffers when the mole (or molarity) ratio of acid to base is between 0.1 and 10.

  8. Therefore, if we want to prepare a buffer that works well near some specified pH, we look for an acid with a Ka as close to the desired pH as possible. • The tool that describes the necessary relationship between pKaand pH to successfully prepare a buffer solution: • pH = pKa ± 1

  9. Henderson Hasselbalch equation: • If you take a biology course, you’re likely to run into a logarithmic form of this equation :

  10. Henderson Hasselbalch equation: • And rearranging to get • pH= pKa + log [A-] • [HA] • It is called the Henderson Hasselbalchequation. The shape of the titration curve of any weak acid is described by the Henderson - Hasselbalch equation, which is important for understanding buffer action and acid-base balance in the blood and tissues of vertebrates. • pH= pKa + log [salt ] • [acid]

  11. Henderson Hasselbalch equation: • This equation also allows us to: • (1) calculate pKa, given pH and the molar ratio of proton donor and acceptor • (2) calculate pH, given pKa and the molar ratio of proton donor and acceptor • (3) calculate the molar ratio of proton donor and acceptor, given pH and pKa.

  12. 1- Calculate the pKa of lactic acid, given that when the concentration of lactic acid is 0.010 M and the concentration of lactate is 0.087 M, the pH is 4.8 Solving Problems Using the Henderson- Hasselbalch Equation

  13. Solving Problems Using the Henderson- Hasselbalch Equation • 2- Calculate the pH of a mixture of 0.10 M acetic acid and 0.20 M sodium acetate. The pKa of acetic acid is 4.76

  14. Solving Problems Using the Henderson HasselbalchEquation • 3- Calculate the ratio of the concentrations of acetate and acetic acid required in a buffer system

  15. Thank you

More Related