The quality of a buffer is determined by its buffer capacity, i.e, its resistance to changes in pH when strong acids or bases are added, the buffer capacity corresponds to the amount of H+ or OH– ions that can be neutralised by the buffer. The effectiveness of three buffering solutions containing acetic acid, Tris and glycine were investigated by titrating them against a weak acid/base and using a pH meter to recored their values. Overall the buffering ranges, where in line with the published pKa values.
A buffer is an aqueous solution consisting of a mixture of a weak acid and its conjugate base or a weak base and 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 changes 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. When some strong acid is added to an equilibrium mixture of the weak acid and its conjugate base, the equilibrium is shifted to the left, in accordance with Le Chatelier's principle. 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. Many life forms thrive only in a relatively small pH range so they utilize a buffer solution to maintain a constant pH. Buffer solutions are necessary to keep the correct pH for enzymes in many organisms to work. When the pH of a buffer solution is one unit above or below its pKa, it will work less effectively.
Materials and Methods:
Titration using a burette and conical flask. Refer to manual for further information.
The graph Figure.1, show the results obtained from the titration of 0.2M NaOH (acid) against 0.1M glycine (base), at 20.6˚c. There is a steady incline in the pH of the solution as more NaOH was added to the glycine. The pH starts off at 7.28 which is the pH of glycine before any HCL is added. The red dot shows the pKa value of the glycine buffer(9.6).
The titration of 0.2M of NaOH against 0.1M of acetic acid, was done at a room temperature 21.2˚c. As seen in Figure.2 the starting pH of the acetic acid, with no NaOH added was 3.22, after adding 0.5 ml of NaOH to the acid this rose to 3.74. Continuing to add between 0.5-0.6 ml of NaOH from the burette at a time, we can see that the pH slowly but gradually increases. There is a slight increase from pH 6.04 to pH 6.88 after which point the pH rapidly increases to 11.04. After this point the pH then continues to steadily increase with the more NaOH that is added.
0.2M of HCL was titrated against 0.1M of Tris, at a temperature of 21.1˚c. Figure.3 shows the a line graph that illustrates the results. 25 ml of Tris was placed into a conical flask, and HCL into a burette, gradually adding between 0.45-0.6 ml of HCL at a time. The pH steadily and gradually decreased. The red dot indicated the point at which the pKa is reached.
Figure.4 displays the results from the titration of 0.2M of NaOH against 0.2M of HCL, at temperature 21.2˚c. We can see that the starting pH of HCL, with no NaOH added is 1.13. This was what was expected given that HCL is an acid., so would have a fairly low pH. We started of by adding amounts of between 1-3 ml of NaOH at a time. Once we began adding the NaOH, the pH somewhat plateaued and there was no real increase or decrease with the increasing amount of NaOH being added to the solution. After the solution gets to a pH of 2.63 and reaches it equivalence point there is a sudden spike in pH up to 10.10 once it reaches it equivalence point, at which time it then continues to steadily increase and then begins to plateau once again when it gets to a pH of 11.60.
Discussion and conclusions:
At the beginning to each experiment the pH