Exp1_Sp2010 - 1-1 Experiment 1 BUFFERS AND pH A large...

Info iconThis preview shows pages 1–3. Sign up to view the full content.

View Full Document Right Arrow Icon
1-1 Experiment 1 BUFFERS AND pH A large fraction of the constituents of cells are weak acids. Proteins, nucleic acids, nucleotides, fatty acids, amino acids and most metabolic intermediates are a few cases in point. At low pH, weak acids are protonated; at high pH, unprotonated. Since the acquisition of a proton can cause an uncharged base to take on a positive charge ( i.e., NH 3 + H + ± NH 4 + ) or can neutralize a negative charge ( i.e., RCOO + H + ± RCOOH), the ionic forms of the many molecules that exist in a cell are very much dependent on the intracellular pH. Biologically important weak acids have a wide range of acid dissociation constants. This creates two challenges for the experimentalist. The first is to determine which ionic forms are most appropriate in a particular cellular/experimental setting. The second is to set and maintain the pH at a value that will assure appropriate levels of the "biologically active" ionic forms of a particular weak acid. This second challenge is also faced by the cell as it conducts its normal functions in an environment where protons are being produced and/or consumed in a myriad of reactions. The pH is set by fixing the ratios of protonated and unprotonated forms of all ionizable groups within the solution. The relation between pH and the ratios of the protonated and unprotonated forms of weak acids and bases is described by the Henderson-Hasselbalch equation. The Henderson-Hasselbalch equation is revisited here because it is not only experimentally important in the design of buffers, but also is central to the understanding of many laboratory procedures; for example, separating and identifying molecules, determining pK a values, moderating chemical reactivities, etc. Objectives 1. Prepare buffers, measure their pHs. 2. Determine the pKa of a pH indicator dye. 3. Examine buffer capacity. Theory The dissociation of a Bronsted (protonic) general acid, HA , can be represented by the equation: H A z ± H + + A z–1 ( 1 ) where z and z-1 are the net charges on HA z , an acid, and A z-1 , its conjugate base. Examples of Bronsted acids are: Bronsted Acid z value NH 4 + +1 CH 3 COOH 0 H 3 N + CH 2 COO 0 HPO 4 –2 -2
Background image of page 1

Info iconThis preview has intentionally blurred sections. Sign up to view the full version.

View Full DocumentRight Arrow Icon
1-2 Note, especially in the case of zwitterionic glycine, that z is the algebraic sum of the charges on that species. The z superscripts have been omitted from equations (2) to (5) for clarity. The Law of Mass Action establishes a quantitative relationship between the chemical activities of an acid and its dissociation products: K a = ( a H + )( a A ± ) a HA ( 2 ) where K a is a constant (at constant temperature and pressure) and each a is the activity of the species indicated by the subscript. Activity is a measure of the reactivity of a chemical species. The activity and the concentration of a species do not generally have the same numerical value, but they approach the same value in very dilute solutions. A relationship parallel to that expressed in equation (2), but involving the concentrations of buffering species instead of their activities, is: ± K a = a H + [ A ] [ HA ] ( 3 ) where [A
Background image of page 2
Image of page 3
This is the end of the preview. Sign up to access the rest of the document.

This note was uploaded on 04/07/2010 for the course MCB 102,103,12 taught by Professor Segel during the Spring '10 term at UC Davis.

Page1 / 13

Exp1_Sp2010 - 1-1 Experiment 1 BUFFERS AND pH A large...

This preview shows document pages 1 - 3. Sign up to view the full document.

View Full Document Right Arrow Icon
Ask a homework question - tutors are online