1. pH measurements

2. Measuring the pH of Solutions The activity of a single ion cannot be measured Only measure our ‘best approximation’ to the hydrogen ion activity. Couple a hydrogen electrode with another reference electrode, e.g., a calomel reference electrode.

3. A Cell for ‘Measuring’ the pH Half-cell reactions. HgCl 2 (s) + 2e -  Hg (l) + 2 Cl - (aq) E  (SCE) = 0.2415 V 2 H + (aq) + 2e -  H 2 (g) E  (H + /H 2 ) = 0.000 V Cell Reaction HgCl 2 (s) + H 2 (g)  Hg (l) + 2 H + (aq) + 2Cl - (aq) E  cell = (0.2415 - 0.0000 V) = 0.2415 V Pt H 2 (g), f=1 H + (aq) HgCl 2 Hg Cl - (aq), 3.5 M Pt

4. The Nernst Equation For the above cell Note – concentration of KCl on one side of the liquid junction is so large, the magnitude of the junction potential should be small!

5. The Practical Problem The activity of the Cl - ion in the cell is not accurately known. We try to place the cell in a reference solution with an accurately known pH (solution I). Next place the solution whose pH we are attempting to measure into the cell (solution II).

6. Assuming that the E LJ and the a(Cl - ) are the same in both cases, Substituting the definition of the pH into the above expression,

7. Standard Solutions Generally, two solutions are used as references. Saturated aqueous solution of sodium hydrogen tartarate, pH = 3.557 at 25  C. 0.0100 mol/kg disodium tetraborate, pH = 9.180 at 25  C.

8. The Glass Electrode The glass electrode has replaced the hydrogen electrode in the operational definition of the pH.

9. Glass Electrodes Measuring the pH – the glass electrode is immersed in the solution of interest. Inner solution – generally a phosphate buffer with a sufficient quantity of Cl - (aq). Silver-silver chloride electrode is sealed within the cell; calomel electrode is used as the reference electrode.

10. Glass Electrodes and pH The potential difference arises across the special glass membrane Equilibrate the hydronium ions inside the membrane with those outside the membrane.