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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 Biological activity (enzyme activity) – 4 5 6 7 pH 8 9 10 Dissociation of Water OH H K H 2O By Convention therefore w 1014 [H2O] = 1 [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 pH is a log scale [H+] pH 10-7 7 10-7 10-6 6 10-8 10-5 5 10-9 10-3 3 10-11 10-11 11 10-3 [OH-] Measurement of pH pH meter and glass electrode – quick – easy – accurate – portable Indicators – titrations phenolphthalein: pink colourless below pH 8.3 methyl orange: red yellow above pH 4.3 Weak acids and strong acids An acid is substance produces H+ in water H2SO4 2H+ + SO42- A base produces OH- and/or accepts H+ NaOH Na+ + OH- A strong acid dissociates completely 1 mole HCl 1 mole H+ + 1 mole Cl1 mole H2SO4 2 mole H+ + 1 mole SO42A weak acid dissociates only partially 1 mole CH3COOH 0.0042 mole H+ + 0.0042 mole CH3COOThe concentration of hydrogen ions [H+] is therefore not always the same as the concentration of the acid Buffers Chemicals which resist pH change – Acetic acid Acetate CH3COOH CH3COO- + H+ – Carbonate Bicarbonate CO32- + H+ HCO3- Amphoteric chemicals e.g. Proteins and amino acids (have both +ve and -ve charged groups on the same molecule) – Buffering range of a buffering chemical is indicated by its pKa pKa is the pH at which the buffering chemical is half dissociated: for HA H+ + Awhen [HA] = [H+] = [A-], then pH = pKa therefore buffering greatest when pH = pKa Buffering capacity is given by the amount of buffering chemical present Carbonate-Bicarbonate Buffering Major buffering in aquatic systems CO2 (g) CO2 (aq) CO2 (aq) + H2O H2CO3 (carbonic acid) Difficult to distinguish between the two forms in water. [H2CO3*] = [CO2] + [H2CO3] H2CO3* is a proxy for “dissolved CO2 plus carbonic acid” "Carbonic acid" dissociates to form bicarbonate H2CO3* HCO3- + H+ pKa = 6.3 Bicarbonate dissociates to form carbonate HCO3- CO32- + H+ pKa = 10.3 Carbonate can also come from the dissolution of carbonate containing minerals: MgCO3, Ca CO3 MgCO3 Mg2+ + CO32CaCO3 + CO2(aq) + H2O Ca2+ + 2 HCO3- Carbonate / bicarbonate system in a particular water depends on its contact with air (CO2) and carbonate minerals. For a closed system with no minerals or CO2 input, the species are: 1.0 HCO3- H2CO3 CO32- 0.8 0.6 0.4 0.2 0 4 5 6 7 pKa 6.3 8 pH 9 10 pKa 10.3 11 12 References Sawyer, McCarty, Parkin(1994) Chemistry for Environmental Engineering Snoeyink, V.L. and Jenkins, D. (1980) Water chemistry, Wiley. Stum, J and Morgan, J.J. (1981) Aquatic Chemistry, Wiley Interscience. Loewenthal, R.E. and Marais, G.V.R (1976) Carbonate Chemistry of Aquatic Systems, Butterworths.