Acid-Base Titrations

CH 229 – General Chemistry 23. Acid-Base Titrations Introduction Acid base reactions are one of the most important chemical reactions there is. A common form of this type of reaction is called titration. Titration is the addition of a solution with a known concentration (the titrant) to a solution with a known volume but an unknown concentration (the analyte), you add the known titrant to the analyte till the reaction reaches neutralization. Titration is important in many ways but one specific way it impacts the real world is through our food. Titration allows food companies to determine the exact amount of salt in a certain food they are selling or to know the concentration of vitamin E (Bell-Young). Knowing these values ensures that the food is safe to eat. Titrations are important to chemistry because they allow you to discover the concentration of unknown solutions leading to new discoveries about how solutions react together. In this experiment we titrated a weak acid KHP with strong NaOH. This allowed us to find the molarity of KHP. We then applied this technique to a weak acetic acid solution and was able to calculqte the pH curve and finds its equivalence and half equivalence point. Balanced Chemical Equations (6 points, this includes that the equations are appropriate given the experimental results, which is also part of the discussion) ● Include any chemical equations as NET IONIC EQUATIONS with phase labels you used in the experiment. Do not use abbreviations for any of the substances, use chemical formulas. This should include at least ● The acid-base net ionic equation for the strong base/weak-acid titration occurring in the standardization 1. ࠵?࠵?࠵?(࠵?) + ࠵?࠵? − (࠵?࠵?) ⇌ ࠵? 2 ࠵?(࠵?) + ࠵?࠵? − (࠵?࠵?) ● The acid-base net ionic equation for the strong base/weak-acid titration that you carried out. 2. ࠵?࠵? 3 ࠵?࠵?࠵?࠵?(࠵?࠵?) + ࠵?࠵? − (࠵?࠵?) ⇌ ࠵?࠵? 3 ࠵?࠵?࠵? − (࠵?࠵?) + ࠵? 2 ࠵?(࠵?) ● The hydrolysis reaction (acid or base dissociation net ionic equation) that is responsible for the initial pH of the weak acid. 3. ࠵?࠵? 3 ࠵?࠵?࠵?࠵?(࠵?࠵?) + ࠵? 2 ࠵?(࠵?) ⇌ ࠵?࠵? 3 ࠵?࠵?࠵? − (࠵?࠵?) + ࠵? 3 ࠵? + (࠵?࠵?) ● The hydrolysis reaction (acid or base dissociation net ionic equation) that is responsible for the equivalence point pH in your acid/base titration. 4. ࠵?࠵? 3 ࠵?࠵?࠵? − (࠵?࠵?) + ࠵? 2 ࠵?(࠵?) ⇌ ࠵?࠵? 3 ࠵?࠵?࠵?࠵?(࠵?࠵?) + ࠵?࠵? − (࠵?࠵?) Results: Data, calculations, and graphs. Tables of Data and Calculation Results (10 points) Table 1: Titration of KHP with NaOH

Trial 1 2 3 Mass of KHP (g) .8805 .8811 .8680 Moles of KHP (mol) .004312 .004314 .004250 Volume NaOH 22.00 21.10 20.70 Molarity of KHP (mol/L) .1959 .2043 .2053 Avg Molarity of NaOH .2018 ± 0.00012 Avg Class Molarity NaOH .208 ± 0.016 Observations: The solution was clear and the KHP was a solid white looking salt that took some time for it to dissolve in solution. The titrant made a slight pink cloud in solution when added and faded until the solution passed the equivalence point which is when the solution went light pink. Table 2: Acetic Acid Titration Acid Acetic Acid Volume of Acid (mL) 10 Initial pH 3.01 Equivalence point Volume NaOH(mL) 12.7 Observations: The solution was a clear solution and when we added the titrant to the solution it turned it pink temporarily until the solution was at the equilibrium point where it turned a light pink and the more NaOH we added the darker the pink got. Table 3: Mandelic, Formic, Lactic and Acetic Acid Titrations Acid Mandelic Formic Lactic Acetic Molarity of Acid .229 ± .012 .220 ± .010 .230 ± .017 .261 ± 0.008 Initial pH 2.6 ± .7 2.7 ± .5 2.6 ± .3 2.8 ± .5 pKa 3.4 ± .6 3.6 ± .3 3.6 ± .3 4.5 ± .4 Ka 4. 0࠵?10 −4 2. 5࠵?10 −4 2. 5࠵?10 −4 3. 2࠵?10 −5 % dissociation 11 ± 10 8.1 ± 2.8 8.3 ± 3.3 3.7 ± 1.9 Calculations (8 pts) Part A