Titration Basics Titration = addition of a measurable volume of a known solution (titrant) to an unknown solution until it is just consumed Use the stoichiometry of the reaction of the known and unknown to calculate the concentration of the unknown solution A pH curve shows the change in pH versus volume of titrant as the titration proceeds pH meter can be used to monitor pH during the titration 2) An acid-base indicator can be used to signal reaching the equivalence point First Derivative Curve Shows where change is greatest

3) Important points: pH increases slowly far from the equivalence point pH changes quickly near the equivalence point The equivalence point of a strong acid—strong base titration = 7.00 4) The titration of a strong base with a strong acid is almost identical

Titration of a Weak Acid with a Strong Base Addition of a strong base to a weak acid forms a Buffer Solution HA + OH- A- + H2O If not enough base has been added to complete the reaction: HA/A- buffer B. Important Points pH increases more rapidly at the start than for a strong acid pH levels off near pKa due to HA/A- buffering effect pH = pKa + log([A-]/[HA]) = pKa + log(1) = pKa (when [A-] = [HA]) Curve is steepest near equivalence point. Equivalence Point > 7.0 Curve is similar to strong acid—strong base after eq. pt. where OH- is major ½ Equivalence Pt. pH = pKa 12.5

Titration of a Weak Base with a Strong Acid Similar problem to the titration of a weak acid with a strong base Determine major species from the stoichiometry Calculate pH from weak acid, buffer, or weak base accordingly Example: Titrate 100 ml of 0.10 M NH3 (Kb = 1.8 x 10-5) with 0.1 M HCl.  

IV. Titrations of Polyprotic Acids and Bases Multiple Inflection Points = Multiple Equivalence Points will be seen The volume required to reach each equivalence point will be the same CO32- + H+ HCO3- Kb1 = KW/Ka2 = 1.8 x 10-4 (pKb1 = 3.74) HCO3- + H+ H2CO3 Kb2 = KW/Ka1 = 2.3 x 10-8 (pKb2 = 7.64) ½ Eq. pt 1 Eq. pt 1 ½ Eq. pt 2 Eq. pt 2 pKa2 = 10.26 pKb1 = 3.74 pKa1 = 6.36 pKb2 = 7.64

Acid-Base Indicators Finding the equivalence point of a titration Use a pH meter Plot pH versus titrant volume Center vertical region = equivalence point Use an Acid-Base Indicator Acid-Base Indicator = molecule that changes color based on pH Choose an indicator that changes color at the equivalence point End Point = when the indicator changes color. If you have chosen the wrong indicator, the end point will be different than the eq. pt. Indicators are often Weak Acids that lose a proton (causing the color change) when [OH-] reaches a certain concentration HIn + OH- In- + H2O

B. We can use the Henderson-Hasselbalch equation on Indicators as well pH = pKa + log([In-]/[HIn]) pH = pKa + log(1/10) for a color change log(1/10) = -1 pH for color change starting in acid is always pKa – 1 for any Indicator For a basic solution titrated with acid, [In-]/[HIn] = 10/1 for color change Log(10/1) = +1, pH for color change will equal pKa + 1 Useful range for a pH Indicator is always pKa +/- 1