A buffer protects against rapids changes in pH when acids or bases are added.  Every living cell is buffered to maintain constant pH and proper cell function.  Consumer products are often buffered to safeguard their activity.  The purpose of this lab activity is to investigate how buffers are made, the pH range in which they are effective, and their buffer capacity.

Background Information

The ability of buffers to resist changes in pH upon the addition of an acid or a base can be traced to their chemical composition.  All buffers contain a mixture of both a weak acid (HA) and its conjugate base (A-), which are related to each other by means of the dissociation reaction shown in Equation 1.  An important feature of the dissociation reaction is the double arrow (⇄), which indicates that the reaction is reversible and that both the weak acid and the conjugate base are present at equilibrium.

Equation 1    

Buffers control pH because the two buffer components (HA and A-) are able to neutralized either acid or base added to the solution. The weak acid component HA reacts with any base added to the solution to give its conjugate base A-.  The conjugate base component A- reacts with any acid added to the solution to form its acid partner HA.  These reaction can be visualized as a cyclic process (see Figure 1 below).  Buffer activity will continue as long as neither component A- or HA is completely consumed or overwhelmed by the amount of strong acid or base.

Properties of Weak Acids and Bases

The properties of weak acids and their conjugate bases determine why buffers behave as they do.  The key difference between a weak acid and a strong acid is that the dissociation of a weak acid is reversible and occurs to only a very limited degree in water.  One familiar weak acid is acetic acid (CH3COOH), which is the main ingredient in vinegar.  A 0.1M solution of acetic acid has a hydronium ion [H3O+] equal to 0.0013M, giving an observe pH of 2.8 to 2.9 (recall the definition and mathematical relationship between [H3O+] and pH: pH = -log[H3O+].)  The observed pH value suggests that only about 1% of the acetic acid molecules are dissociated  to the conjugate basse form, acetate ion (CH3COO-) under these conditions. In contrast, a strong acid such as hydrochloric acid (HCl) undergoes complete and irreversible 100% dissociation in water.

The degree to which a weak acid is ionized in aqueous solution is governed by the equilibrium constant Ka for its reversible dissociation reaction (Equation 2 & 3).  The equilibrium constant Ka is also referred to as the dissociation constant of the weak acid.  The Ka value  for acetic acid, for example, is 1.76 x 10-5.

The Buffer Equation

Generalization of Equation 3 for any weak acid HA and its conjugate base A- gives Equation 4, which can, in turn, be rearranged to solve for the [H3O+] concentration (Equation 5).  Equation 5 is sometimes known as the buffer equation;  it provides the key to calculating the properties of buffer solutions. 
When the concentrations of the weak acid  and its conjugate base are equal, the ratio in Equation 5 will be equal to one and the 
[H3O+] concentration will be equal to the dissociation constant Ka for the weak acid.  Careful selection of the identity of the weak acid  component makes it possible to prepare a buffer solution with almost any initial  pH value.  In the case of acetic acid, for example, a buffer solution  consisting of a 1:1 molar mixture of acetic acid and its conjugate base sodium acetate will have a hydronium ion concentration equal to 1.76 x 10-5 M, and the pH of the solution will be 4.75.  Carbonic acid (H2CO3) has a Kvalue equal to 4.3 x 10-7.  A buffer prepared from equal moles of carbonic acid and its conjugate base  bicarbonate ion (HCO3-) will have an [H3O+] concentration equal to 4.4 x 10-7 M and a pH value equal to 6.4.

What happens when strong acid or base is added  to a buffer?  Reaction of the weak acid component HA with additional base , such as sodium hydroxide, convert the weak acid to its conjugate base from A- (Equation 6).  Similarly, reaction of the basic component A- with added acid results in its neutralization to the conjugate acid form HA (Equation 7).
The effect of adding a strong acid or base on the pH of a buffer solution can be predicted using LeChâtelier's principle.  Consider the equimolar acetic acid - acetate buffer (Equation 2).  Adding HCl to the buffer solution, with its equilibrium pH = 4.75, increases the concentration of H3O+ ions, one of the products of the reversible reaction. This shifts the equilibrium to the left, increasing the concentration of acetic acid  and decreasing the concentration of acetate ions.  The ratio of [HA] to [A-] in Equation 5 increases as well increases as well, and [H3O+] is larger - the pH decreases.  The opposite effect is observed when NaOH is added to the buffer solution. OH- ions neutralize some of the  H3O+ ions, which shifts the equilibrium to the right, increasing the concentration of acetate ions relative to acetic acid molecules.  The ratio of [HA] to [A-] decreases , and [H3O+] is smaller - the pH increases.  In either cases, however, as long the [HA]/[A-] ratio stays within certain limits, the pH remains relatively constant. 

Buffer Range and Buffer Capacity

A buffer composed of an equal number of moles of a weak acid and its conjugate base is sometimes called an ideal buffer because it is equally effective in resisting  pH changes upon addition of either acid or base.  As shown in the example above, in an ideal buffer solution the [H3O+] concentration is equal to the dissociation constant (Ka) for the weak acid.

The pH range in which a buffer solution will be effective is called the buffer range.  Since a buffer solution must always contain noticable amounts of both a weak acid and its conjugate base, the buffer range is usually limited to concentration ratios of HA:A- between 1:10 and 10:1.  Substituting these concentration ratios in Equation 8 reveals that the effective pH range for a given buffer is plus or minus one unit on either side of the pH value of the ideal buffer.  An ideal acetic acid-sodium acetate buffer system  has a pH of 4.75 and its buffer range is 3.75-5.75.  Equation 8 shows the calculation for the lower pH limit of an acetic acid-sodium acetate buffer solution (when the concentration ratio of the weak acid component to the conjugate base component is equal to 10:1)   

The effectiveness of a buffer in resisting pH changes is called the buffer capacity. Consideration of Equation 5 reveals that the pH of a buffer prepared form a weak acid HA and its conjugate base A- should be independent of their total concentration as long as the ratio [HA] to [A-] is the same .  Thus, an acetic acid-acetate buffer prepared from 0.1 mole HA and 0.1 mole A- should have the same theoretical pH as a buffer containing 1 mole HA and 1 mole A-.  The buffer capacity of the buffers, however, will be very different.  The capacity of the 0.1 moles HA/0.1 moles A- buffer will be overwhelmed  when approximately 0.09 moles of HCl or NaOH have been added.  The 1 M buffer will withstand almost 10x as much strong acid or strong base before either HA or A- is consumed.  

Experiment Overview

The purpose of this lab is to design and make an effective buffer with a specific pH value for a consumer or experimental biochemistry application. The investigation begins with an introductory activity  to compare the properties of three acetate buffers containing varying ratios of HA and A-.  The results provide a model for desinging an experiment to prepare a desired buffer and verify its properties and performance.  

Five different buffer "challenges" are presented - each student group will be assigned one. The specification for each buffer challenge are that a) the pH should  be within ±0.5 pH units  of the desired, and b) 25 mL of the buffer should maintain the desired pH ±1 after 10 mL of 0.02M HCl or 10 mL 0.2M NaOH have been added. Preparation of a buffer by partial neutralization of a weak acid or a weak base offers additional opportunities for inquiry.