2. Principles of buffering.
a. Simple buffering agents.
b. "Universal" buffer mixtures.
c. Common buffer compounds used in biology.
4. Buffer capacity.
5. Calculating buffer pH
a. Monoprotic acids.
b. Polyprotic acids.
A buffer is an aqueous solution consisting of a mixture of a weak acid and its conjugate base or a weak baseand its conjugate acid. Its pH changes very little when a small amount of strong acid or base is added to it and thus it is used to prevent any change in the pH of a solution. Buffer solutions are used as a means of keeping pH at a nearly constant value in a wide variety of chemical applications. Many life forms thrive only in a relatively small pH range so they utilize a buffer solution to maintain a constant pH. One example of a buffer solution found in nature is blood. So far in discussing pH we have dealt only with solutions obtained by adding a single acid, such as acetic acid, or a single base, such as the acetate ion, to water. We must now turn to a consideration of solutions to which both an acid and a base have been added. The simplest case of such a solution occurs when the acid and base are conjugate to each other and also present in comparable amounts. Solutions of this special kind are called buffer solutionsbecause, as we shall shortly see, it is difficult to change their pH even when an appreciable amount of strong acid orstrong base is added. As a typical example of a buffer solution, let us consider the solution obtained when 3.00 mol acetic acid (HC2H3O2) and 2.00 mol sodium acetate (Na C2H3O2) are added to sufficient water to produce a solution of total volume 1 dm³. The stoichiometric concentration of acetic acid, namely, ca, is then 3.00 mol dm–3, while the stoichiometric concentration of sodium acetate, cb, is 2.00 mol dm–3. As a result of mixing the two components, some of the acetic acid, say x mol dm–3, is converted to acetate ion and hydronium ion.
Principles of buffering
Solution of a weak acid (pKa = 4.7) with Simulated titration of an acidified alkali. [pic]
Addition of hydroxide to a mixture of a weak acid and its conjugate base
Buffer solutions achieve their resistance to pH change because of the presence of an equilibrium between the acid HA and its conjugate base A-.
HA [pic] H+ + A- When some strong acid is added to an equilibrium mixture of theweak acid and its conjugate base, the equilibrium is shifted to the left,
in accordance with Le Chatelier's principle.
Because of this, the hydrogen ion concentration increases by less than the amount expected for the quantity of strong acid added.
Similarly, if strong alkali is added to the mixture the hydrogen ion concentration decreases by less than the amount expected for the quantity of alkali added.
The effect is illustrated by the simulated titration of a weak acid with pKa = 4.7.
The relative concentration of undissociated acid is shown in blue and of its conjugate base in red.
The pH changes relatively slowly in the buffer region, pH = pKa ± 1,
centered at pH = 4.7 where [HA] = [A-]. The hydrogen ion concentration decreases by less than the amount expected because most of the added hydroxide ion is consumed in the reaction.
OH- + HA → H2O + A-
and only a little is consumed in the neutralization reaction which results in an increase in pH.
OH- + H+ → H2O
Once the acid is more than 95% deprotonated the pH rises rapidly
because most of the added alkali is consumed in the neutralization reaction....