1. pH and Buffering  Aim  to know the logarithmic scale of pH  to understand how weakly dissociating acids can buffer the pH of an aqueous environment  to know the importance of the carbonate - bicarbonate buffering system

2. pH, The master variable – Consumed and produced – Enzyme/biological optima 4 5 6 7 8 9 10 pH Biological activity (enzyme activity)

3. By Convention[H 2 O] = 1 therefore [OH - ] [H + ] = 10 -14 So, if [H + ] is known, [OH - ] is also known if [H + ] = 10 -5, then [OH - ] =10 -9 Dealing in [H + ] is cumbersome Deal in pH (minus the log of the hydrogen ion concentration) pH = - log[H + ] if [H + ] = 0.1 M or 10 -1 M, then pH = 1 Dissociation of Water

4. pH is a log scale pH 10 -7 7 10 -6 6 10 -8 10 -5 510 -9 10 -3 310 -11 1110 -3 [H + ][OH - ]

5. n pH meter and glass electrode – quick – easy – accurate – portable n Indicators – titrations phenolphthalein: pink  colourless below pH 8.3 methyl orange: red  yellow above pH 4.3 Measurement of pH

6. n An acid is substance produces H + in water H 2 SO 4  2H + + SO 4 2- n A base produces OH - and/or accepts H + NaOH  Na + + OH - n A strong acid dissociates completely 1 mole HCl  1 mole H + + 1 mole Cl - 1 mole H 2 SO 4  2 mole H + + 1 mole SO 4 2- n A weak acid dissociates only partially 1 mole CH 3 COOH  0.0042 mole H + + 0.0042 mole CH 3 COO - n The concentration of hydrogen ions [H + ] is therefore not always the same as the concentration of the acid Weak acids and strong acids

7. Buffers n Chemicals which resist pH change – Acetic acid Acetate CH 3 COOH  CH 3 COO - + H + – Carbonate Bicarbonate CO 3 2- + H +  HCO 3 - n Amphoteric chemicals – e.g. Proteins and amino acids (have both +ve and -ve charged groups on the same molecule)

8. n Buffering range of a buffering chemical is indicated by its pK a pKa is the pH at which the buffering chemical is half dissociated: for HA  H + + A - when [HA] = [H + ] = [A - ], then pH = pKa therefore buffering greatest when pH = pKa n Buffering capacity is given by the amount of buffering chemical present

9. Major buffering in aquatic systems CO 2 (g)  CO 2 (aq) CO 2 (aq) + H 2 O  H 2 CO 3 (carbonic acid) Difficult to distinguish between the two forms in water. [H 2 CO 3 *] = [CO 2 ] + [H 2 CO 3 ] H 2 CO 3 * is a proxy for “dissolved CO 2 plus carbonic acid” Carbonate-Bicarbonate Buffering

10. "Carbonic acid" dissociates to form bicarbonate H 2 CO 3 *  HCO 3 - + H + pK a = 6.3 Bicarbonate dissociates to form carbonate HCO 3 -  CO 3 2- + H + pK a = 10.3 Carbonate can also come from the dissolution of carbonate containing minerals: MgCO 3, Ca CO 3 MgCO 3  Mg 2+ + CO 3 2- CaCO 3 + CO 2 (aq) + H 2 O  Ca 2+ + 2 HCO 3 -

11. Carbonate / bicarbonate system in a particular water depends on its contact with air (CO 2 ) and carbonate minerals. For a closed system with no minerals or CO 2 input, the species are: 1.0 0.8 0.6 0.4 0.2 0 Fraction as designated species H 2 CO 3 HCO 3 - CO 3 2- 4 57610891211 pH pK a 6.3 pK a 10.3

12. References n Sawyer, McCarty, Parkin(1994) Chemistry for Environmental Engineering n Snoeyink, V.L. and Jenkins, D. (1980) Water chemistry, Wiley. n Stum, J and Morgan, J.J. (1981) Aquatic Chemistry, Wiley Interscience. n Loewenthal, R.E. and Marais, G.V.R (1976) Carbonate Chemistry of Aquatic Systems, Butterworths.