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

pH, The master variable – Consumed and produced – Enzyme/biological optima pH Biological activity (enzyme activity)

By Convention[H 2 O] = 1 therefore [OH - ] [H + ] = 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 M, then pH = 1 Dissociation of Water

pH is a log scale pH [H + ][OH - ]

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

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 mole Cl - 1 mole H 2 SO 4  2 mole H mole SO 4 2- n A weak acid dissociates only partially 1 mole CH 3 COOH  mole H 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

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

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

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

"Carbonic acid" dissociates to form bicarbonate H 2 CO 3 *  HCO H + pK a = 6.3 Bicarbonate dissociates to form carbonate HCO 3 -  CO 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 HCO 3 -

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: Fraction as designated species H 2 CO 3 HCO 3 - CO pH pK a 6.3 pK a 10.3

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.