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Understanding pH and Buffering Systems in Aqueous Environments

Explore the logarithmic pH scale, weakly dissociating acids, importance of buffering systems, and practical aspects of pH measurement using indicators and meters. Learn about strong and weak acids, chemicals resisting pH change, buffers, and the carbonate-bicarbonate buffering system.

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Understanding pH and Buffering Systems in Aqueous Environments

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  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 Biological activity (enzyme activity) 4 5 6 7 8 9 10 pH

  3. Dissociation of Water By Convention [H2O] = 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

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

  5. 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

  6. 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 Cl- 1 mole H2SO4 2 mole H+ + 1 mole SO42- • A weak acid dissociates only partially 1 mole CH3COOH  0.0042 mole H+ + 0.0042 mole CH3COO- • The concentration of hydrogen ions [H+] is therefore not always the same as the concentration of the acid

  7. 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) Buffers

  8. 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+ + A- when [HA] = [H+] = [A-], then pH = pKa therefore buffering greatest when pH = pKa • Buffering capacity is given by the amount of buffering chemical present

  9. 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”

  10. "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+ + CO32- CaCO3 + CO2(aq) + H2O  Ca2+ + 2 HCO3-

  11. 1.0 0.8 0.6 0.4 0.2 0 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: HCO3- H2CO3 CO32- Fraction as designated species 5 6 7 8 9 10 11 12 4 pH pKa 10.3 pKa 6.3

  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.

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