Chapter 7 Mixtures of Acids and Bases Common-Ion Effect Common-Ion Effect: Shift in position of an equilibrium caused by addition of an ion involved in the reaction. Essentially the addition of a conjugate acid or base in the form of the appropriate salt. Example: pH Buffers Buffer: Henderson-Hasselbalch Equation Sample Exercise - Basic Buffer An Environmental Buffer A Physiological Buffer Buffer Range and Capacity pH Change of Buffer vs Buffer Concentration Sample Exercise - Buffering Action pH Indicators pH Titrations Titration Curves for NH3 and NaOH Titration Curves for Acetic and Hydrochloric Acids Detailed Sample Exercise on Strong Acid/Weak Base Titration Suppose 50.00 mL of pH 10.0 water from a hot spring rich in carbonate and bicarbonate ions, is titrated with 0.02075 M HCl. A few drops of phenophthalein are added at the beginning of the titration, and the solution turns pink. It takes 11.21 mL of titrant to reach the pink-to-clear equivalence point. Then a few drops of bromocresol green are added, and it takes an additional 32.28 mL of titrant before the blue-green color changes to yellow. What are the initial concentrations of carbonate and bicarbonate in the sample? We are asked to determine the concentrations of two analytes in one sample from the results of a single titration with a monoprotic strong acid, HCl. These determinations are based on the volumes of titrant needed to reach two equivalence points: an initial one at which any CO32− in the sample has been converted to HCO3−: (1) H+(aq) + CO32−(aq) → HCO3−(aq) and a second one at which HCO3− has been converted to H2CO3 (2) H+(aq) + HCO3-(aq) → H2CO3(aq) The HCO3- titrated in reaction 2 includes any HCO3- in the original sample plus all the HCO3- produced in reaction 1. If there were no HCO3- present initially, the volume of titrant needed to react with the HCO3- in reaction 2 would be exactly the same as the volume needed to react with CO32-, 11.21 mL, in reaction 1. However, the volume of the titrant required to reach the second equivalence point is much greater: 32.28 mL. The difference between these two volumes, (32.28 - 11.21) = 21.07 mL, is the volume of acid required to react with any HCO3- in the sample, the value of [HCO3-] that we calculate for the original sample shouldbe nearly twice its [CO3-] value.  Alkalinity Titration: CO3- Ion (1) CO3-(aq) + H+(aq) ⇌ HCO3-(aq) (2) HCO3-(aq) + H+(aq) ⇌ H2CO3(aq) Sample Exercise (cont.) The stoichiometry of the reaction tells us that the titrant and carbonate react in a 1:1 mole ratio, so, at the first equivalence point, mol HCl added = mol CO32- consumed This means that the ratio nA/nB for this reaction is 1:1. The same is true for the conversion of HCO3- into H2CO3: mol HCl added = mol HCO3- consumed Using VAMA = (nA/nB)VBMB- Note that nA/nB is 1/1 in this example The titration results confirm that the bicarbonate concentration in the original sample was almost twice the carbonate concentration. To check our two values from the calculations, we can insert the results into the Henderson-Hasselbalch equation and calculate what the initial pH of the sample should have been. To do that, we need the value for carbonic acid, which is 10.33. This calculated pH value agrees with the pH given in the problem