Then plugging into the pH equation:
The pH has shifted from the neutral pH of 7 for the pure water to a pH of almost 12 with the NaOH added, a shift of 5 pH units!
Buffering agents stabilize the pH of aqueous solutions
In example 5, the pH of the buffer solution changed by less than 0.5 units when the NaOH was added. Yet adding the same amount of NaOH to pure water resulted in a 5 unit pH shift in example 6. Clearly, the conjugate acid/base pair was able to protect the solution from a large change in pH.
A buffer is a solution that resists a change in pH when acids or bases are added to it.
How is this possible? Buffers work by acting a little bit like a sponge, soaking up excess H3O+ or OH ions when they are added to a solution. The dissociation reaction for acetic acid shows that a solution of acetic acid contains acetate ions:
How do acetic acid and the acetate ion work to buffer the solution? If a strong acid, such as hydrochloric acid, is added to this buffer, the H3O+ ions generated will react with the acetate ions, removing them from the solution:
Similarly, if a strong base, generating lots of OH ions, were added to the acetic acid buffer solution
The OH ions are removed via a reaction with acetic acid. It is important to remember that buffers cannot maintain their pH indefinitely as more acid or base is added. Imagine slowly pouring a bucketful of water onto a small sponge on the kitchen floor. At first, as the water pours out of the bucket it is absorbed by the sponge, keeping the kitchen floor dry. But once the sponge is soaked through, it can hold no more and the water spills onto the floor as fast as if the sponge werent there. In the same way, buffers will protect the pH of the solution to some extent, but if they are inundated by large amounts of H3O+ or OH ions, all the available conjugate acid or base molecules will have been used up and pH changes will rapidly occur.
Thus, buffers only work as long as the amount of conjugate acid and base ions are large compared to the H3O+amount of or OH ions to be removed. Remember from the Henderson-Hasselbalch equation (section 5) that the numbers of conjugate acid molecules is equal to the number of conjugate base molecules when the pH = pKa:
Because the conjugate acid and conjugate base molecules are available in equal amounts when the pH = pKa, the buffer solution has the strongest ability to protect against pH changes caused by incoming H3O+ or OH ions when the pH of the solution is close to the pKa of the conjugate acid. If the pH of the solution strays too far from the pKa point, the buffer uses up all its available conjugate acid or base molecules, and the pH starts to fall (or rise) dramatically. At this point, the buffering capacity of the buffer has been exceeded. this can be seen in the buffer (weak acid) titration graph below.
Buffers protect the pH of a solution best within one pH unit of the pKa.
Buffers are important for biochemistry because the structures (and therefore the functions) of biological molecules are stable within a relatively narrow range of pH values.The evolution of life, which is believed to have begun in water, was likely due in part to the stable pH of seawater. Even today, the largest buffered systems in the world are the Earth's oceans. Carbon dioxide (CO2) from the Earths atmosphere reacts with H2O to produce carbonic acid, H2CO3. Carbonic acid in turn reacts with water to form bicarbonate:
Furthermore, the bicarbonate ion can also react with water to form carbonate:
The carbonate and bicarbonate ions act together as conjugate acid/base pair to keep the pH of the ocean at about 8.2.
In the living organisms, proteins (the molecules which do most of the work of a cell) are very sensitive to acid concentration, because they only fold into their proper three-dimensional shapes within a small pH range. Some of the individual amino acid building blocks that make up a protein have ionizable side chains that change their ionization state, or charge, at different pH. Thus, if the pH changes, some of the attractive forces that hold the protein together will also change.
The ionization states of the side chain of the amino acid lysine are shown below. The pKa of lysine's side chain amino group is approximately 10.8. Therefore, at a pH lower than the pKa, the side chain is in the protonated (conjugate acid) state; whereas at a pH higher than the pKa, the side chain is unprotonated (conjugate base).
Proteins consist of long chains of amino acids, many of which include acidic or basic groups. These groups are chaged at physiological pH, and form ionic interactions with each other that contribute to the stability of the folded protein. A change in pH of the protein's aqueous environment may alter the charge of some of these charged groups, disrupting the ionic interactions and the stability of the proteins structure.
Remember that the Henderson–Hasselbalch equation was used for calculating the pH of a solution containing a weak acid and its conjugate base. Since such a solution is a buffer, the Henderson–Hasselbalch equation is extremely useful for calculations involving buffers, allowing the following:
pH = pKa + log