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Unformatted text preview: 1-1Experiment 1WEAK ACID BUFFERSIntroductionA large fraction of the constituents of cells are weak acids, and some are weak bases; for example: proteins and individual amino acids, nucleic acids and individual nucleotides, fatty acids, and most metabolic intermediates. Since the acquisition of a proton can cause an uncharged-base to take on a positive charge (i.e. NH3+ H+↔ NH4+) or can neutralize a negative charge (ie RCOO–+ H+↔ RCOOH), the ionic forms of the many molecules that exist in a cell are very much dependent on the intracellular pH. For experiments in vitro, the pH must be set and maintained at a value that will assure appropriate levels of the "biologically active" ionic form(s) of the molecule(s) being examined. The maintenance of pH is accomplished by the introduction of a buffer into the biochemical solution. Buffers resist changes in pH 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 pKa values, moderating chemical reactivity, etc.Objectives1. To prepare buffers, measure their pHs, and examine their buffer capacities. 2. To determine the pKa of a pH indicator dye spectrophotometrically.TheoryThe dissociation of a Bronsted (protonic) general acid, HA, can be represented by the chemicalequation: .HAZ↔ H++ AZ-1(1)where zand z-lare the net charges on HAZ, an acid, and, AZ-1, its conjugate base. Examples of Bronsted acids are: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: Ka = (aH+) (aA–)= ([H+] γH+) ([A–] γA–)(2)aHA[HA] γHABronsted Acidz valueNH4++1CH3COOHH3N+CH2COO-HPO4-2-21-2where Kais a constant (at constant temperature and pressure) and ais the activity of the species. Activity is a measure of the reactivity of a species and can be equated to the concentration multiplied by the activity coefficient, γ, of the species. Under ideal conditions of dilute solutions the activity coefficient is ~1.0, thus the activity is equated to the concentration of a species. Substituting in concentrations of A and HA, and rearranging the equation results in the following expression:Ka'=aH+[A–][HA](3)where Ka'is the apparent equilibrium constantand Ka' is a function of the various species present in the buffer solution (ionic strength, interactions between species), and possibly by temperature.the buffer solution (ionic strength, interactions between species), and possibly by temperature....
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This note was uploaded on 11/21/2010 for the course MCB 120L 69059 taught by Professor Fairclough during the Fall '10 term at UC Davis.
- Fall '10