by John Gipps
John Gipps is a professional officer in the Faculty of Education at Monash University and is involved in the training of primary teachers and secondary science teachers.
Buffers are mixtures that resist changes in pH when acids or alkalis are added. They play a crucial role in regulating chemical conditions in animals andplants and also in global environmental systems such as the oceans. In this article I use computer-interfaced titrations to demonstrate the behaviours of some important buffer systems. consider to be relevant to the concept buffer is a solution of a of buffers. substance or mixture that The Victorian VCE Chemistry Unit resists pH changes when acids 3 topic 'Equilibrium' covers the pH or bases areadded to it. The typical of solutions of weak acids and the buffer is a mixture of a weak acid HA applications of chemical equilibrium in living systems (page 23 of study and its conjugate base A. guide), the VCE Biology Unit 2 topic (1) HA + H.O ^ A- + up'Functioning Organisms' covers (2} A + H,O
The carbonate system shows two buffering ranges, one of 5.5 to 7.5 and the other 9 to ILThe HCO^/ H^CO,/ CO, mixture plays the major role in buffering the pH of blood to a narrow region around 7.4 (Staub, 1991), although hemoglobin is also a significant contributor. A pH of 7.4 is
50 75 0.1 MHCImL
Figure 4 Carbonate Buffer
responsible for keeping it in this range (Home 1969). A quick glance at the titration curve in Figure 4 would suggest that the carbonatesystem ought to be almost useless as a buffer at pH 8.2, but seawater also contains calcium ions, which complicate the picture. For the final experiment I put 0.01 M NaOH in the burette and 50 mL of 0.01 M NaHCO3 mixed with 50 mL of 0.05 M CaClj in the beaker. Figure 5 shows the titration curve for this mixture. The pH initially increased from 7.9 to
very close to the alkaline end of thebuffering range and the system works only because the substances most likely Carbonic acid has two ionisation to enter the blood stream and affect the stages. pH are adds such as carbon dioxide HpO, + H p HCO; + H30*pKa = 6,4 and lactic acid. Regulation of breathing HCO, + H p ^- CO3' + H3O^ pKa - 10.3 is closely connected to the maintenance Carbonic acid itself usually forms only of this pH of 7.4. Ifwe hyperventilate a small part of such mixtures as much so that carbon dioxide is removed from the blood stream faster than metaboUc of it decomposes to carbon dioxide. processes can replace it, the pH will HpO, ^ CO,(aq) + H p become too high. If, on the other hand, CO>q) CO^(g) we have a condition such as asthma
TEACHING SCIENCE I VOLUME 51 NO 1 I AUTUMN 2005
8.3, thenactually decreased to 8.0 even though hydroxide was still being added. At about the same time as this decrease the clear solution became cloudy as calcium carbonate precipitated. HCO,- + OH- -^ CO/- + H,O Ca=^ + CO^^ -- CaCO,(s) The calcium carbonate did not precipitate instantaneously, and in the small delay the levels of carbonate built up. When the precipitation occurred much of the carbonate wasremoved from solution, leaving a drop before there is significant damage to marine ecosystems.
Chester, R. (2000). Marine geochemistry. Maiden, United States: Blackwell Appendix Publishing. Solutions Commons, C, Jarrett, S., McKenzie, Hydrochloric acid 0.1 M was prepared C, Mosely, W., Porter, M. and by diluting 10 mL of concentrated HCl Williamson, M. (1999). Chemistry Two to 1000mL and standardising against (3"^ Ed.). Melbourne: Heinemann. dried potassium carbonate. Cipps, J. (1994). Acids, bases and Sodium hydroxide 0.1 M was prepared computers. Australian Science Teachers by dissolving 1.00 g of NaOH in 250 journal, 40(2), 48-52. mL of water and standardising against Home, R.A. (1969). Marine chemistry. the 0.1 M HCl. New York: Wiley Interscience. James, M., Stokes,...