{[ promptMessage ]}

Bookmark it

{[ promptMessage ]}

D_+Experiment+01+_v+F10 - 1-1 Experiment 1 WEAK ACID...

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

View Full Document Right Arrow Icon
1-1 Experiment 1 WEAK ACID BUFFERS Introduction A 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. NH 3 + H + ↔ NH 4 + ) 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. Objectives 1. To prepare buffers, measure their pHs, and examine their buffer capacities. 2. To determine the pKa of a pH indicator dye spectrophotometrically. Theory The dissociation of a Bronsted (protonic) general acid, HA , can be represented by the chemical equation: . HA Z ↔ H + + A Z-1 (1) where z and z-l are the net charges on HA Z , an acid, and, A Z-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 = (a H+ ) (a A– ) = ([H + ] γ H+ ) ([A ] γ A– ) (2) a HA [HA] γ HA 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 Document Right Arrow Icon
1-2 where Ka is a constant (at constant temperature and pressure) and a is 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: K a ' = a H + [ A ] [ HA ] (3) where K a ' is the apparent equilibrium constant and K a ' is a function of the various species present in the buffer solution (ionic strength, interactions between species), and possibly by temperature.
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.

{[ snackBarMessage ]}