Potentiometric Analysis - Lab Report

Potentiometric Analysis Abstract: The purpose of this lab was to determine the molar mass, molar concentration, stoichiometric point, and pKa of the unknown weak acid, “Ron”, using the graphs created from the data generated by each titration trial. It was hypothesized that the titration curves resulting from the relationship between the addition of NaOH and the pH of the solution would provide enough information to determine each of the desired values. This was supported by the results of the experiment. The graphs revealed the stoichiometric point visually, which allowed the pKa to be found (6.0). It also allowed the molar concentration of 0.620M to be found, as well as the molar mass of 163.6 g/mol. Introduction: The purpose of this experiment was to utilize a pH meter for the formation of a graph from which the stoichiometric point of a weak acid could be determined. Another objective was to determine the molar concentration, molar mass, and pKa value of a weak acid. These objectives were accomplished using titration as a technique to gather the data that was needed to calculate the desired values. A potentiometer was also used to detect ionic concentration differences in the solution. The specific potentiometer used was a pH meter, which showed the pH of the solution as the standardized NaOH solution was added. The volume of NaOH dispensed during each titration made it possible to then calculate the molar concentration as well as the molar mass of the weak acid. The pH meter in conjunction with the titration measurements allowed titration curves to be created for each trial that was performed. The graphs provided the information that was needed to determine the pKa value of the unknown acid. Since the acid that was used was monoprotic, there was only one pKa for each trial of the experiment. It was originally hypothesized that the titration curves showing the relationship between volume of NaOH added and the pH of the solution would provide sufficient information to determine the pKa of the unknown acid called “Ron” in this experiment. Materials and Methods Please refer too Experiment 18, Potentiometric Analyses, on pages 231-233 in Laboratory Manual for Principles of General Chemistry, 10th Edition, by J.A. Beran. A deviation from the published procedure was that only two trials were completed rather than 3. Results: A. Molar Concentration of a Weak Acid Trial 1 1. Molar concentration of NaOH (mol/L) Trial 2 0.100 2. Volume of weak acid (mL) 25 25 3. Buret reading of NaOH, initial (mL) 0 0 4. Buret reading of NaOH at stoichiometric point, final (mL) 15.43 16.11 5. Volume of NaOH dispensed (mL) 15.43 16.11 7. Moles of NaOH to stoichiometric point (mol) 0.015 0.016 8. Moles of acid (mol) 0.015 0.016 9.Molar concentration of acid (mol/L) 0.600 0.640 10. Average molar concentration of acid (mol/L) 0.620 Table 1: Molar Concentration of a Weak Acid This table depicts data and calculations necessary to determine the molar mass of the weak acid for Part A of the experiment. B. Molar Mass and the pKa of a Solid Weak Acid Trial 1 1. Mass of dry, solid acid (g) 2. Molar concentration of NaOH (mol/L) 3. Buret reading of NaOH, initial (mL) 4. Buret reading of NaOH at Trial 2 0.306 0.294 0 0 18.96 17.71 0.100 stoichiometric point, final (mL) 5. Volume of NaOH dispensed, (mL) 18.96 17.71 7. Moles of NaOH to stoichiometric point, (mol) 0.001896 0.001771 8. Moles of acid (mol) 0.001896 0.001771 161.4 166.0 5.54 6.20 6 6 9. Molar mass of acid (g/mol) 10. Average molar mass of acid (g/mol) 163.7 11. Volume of NaOH halfway to stoichiometric point (mL) 12. pKa1 of weak acid (from graph) 13. Average pKa1 6 Table 2: Molar Mass and the pKa of a Solid Weak Acid This table depicts the data and calculations necessary to determine the molar mass and pKa of the solid weak acid used in Part B of the experiment. Figure 1 shows the titration curve that illustrates the relationship between the volume of NaOH added and the pH of the solution for Trial 1 of Part A of the experiment. Figure 2 shows the titration curve that illustrates the relationship between the volume of NaOH added and the pH of the solution for Trial 2 of Part A of the experiment. rve that ween the e pH of of the urve that tween the he pH of t B of the Calculations: A.5. Volume of NaOH dispensed (mL) Final Buret reading (mL) – Initial Buret reading (mL) Trial 1: 15.43 – 0.00 = 15.43 mL Trial 2: 16.11 – 0.00 = 16.11 mL A.7. Moles of NaOH to stoichiometric point (mol) [(Molar concentration of NaOH (mol/L))*(Volume of NaOH dispensed (mL))] / 100 Trial 1: (0.100*15.43) / 100 = 0.015 mol Trial 2: (0.100*16.11) / 100 = 0.016 mol A.9. Molar concentration of acid (mol/L) Moles of acid (mol) / Volume of weak acid (L) Trial 1: 0.015 / 0.025 = 0.600 mol/L Trial 2: 0.016 / 0.025 = 0.640 mol/L A.10. Average molar concentration of acid (mol/L) (Molar concentration of acid in Trial 1 + Molar concentration of acid in Trial 2) / 2 (0.600 + 0.640) / 2 = 0.620 mol/L B.5. Volume of NaOH dispensed (mL) Final buret reading (mL) – Initial buret reading (mL) Trial 1: 18.96 – 0 = 18.96 mL Trial 2: 17.71 – 0 = 17.71 mL B.7. Moles of NaOH to stoichiometric point (mol) [(Molar concentration of NaOH (mol/L))*(Volume of NaOH dispensed (mL)]/1000 Trial 1: (0.100*18.96)/1000 = 0.001896 mol Trial 2: (0.100*17.71)/1000 = 0.001771 mol B.9. Molar mass of acid (g/mol) (Mass of dry, solid acid (g)) / Moles of acid (mol) Trial 1: 0.306 / 0.001896 = 161.4 g/mol Trial 2: 0.294 / 0.001771 = 166.0 g/mol B.10. Average molar mass of acid (g/mol) (Molar mass of acid from Trial 1 + Molar mass of acid from Trial 2) / 2 (161.4 + 166.0) /2 = 163.7 g/mol Discussion: The main purpose of this lab was to use the relationship between volume of NaOH titrated and pH of the solution to graphically determine the stoichiometric point, molar concentration, molar mass, and pKa value of an unknown weak acid. For each titration, the pH meter recorded the changing pH of the acid-base solution while the volume of sodium hydroxide solution being added was also being recorded. This provided the data needed to create the titration curves. As seen in Figures 1-4, there was a gradual increase in pH as NaOH was added at first followed by a steep increase once the volume dispensed got to about 11 mL. The middle of this steep increase was the endpoint, also known as the stoichiometric point. At this point, the amount of base added to the solution was able to neutralize the unknown weak acid, “Ron”. This point was found by measuring the distance between the two intersection points of the tangent lines and marking exactly halfway between the two points. Each figure also shows a point that was marked halfway to the stoichiometric point. At this point, pH = pKa. Therefore, using the graph to determine the pH of the solution at this point also determined the pKa for each trial. The average pKa from Part B of the experiment was found to be 6. Since the moles of sodium hydroxide are equal to the moles of the unknown acid at this point as well, it made it possible to also calculate its molar mass, which was found to be 163.7 g/mol The molar concentration of “Ron” was calculated using the information provided by the titration process in Part A of the experiment. The average molar concentration was found to be 0.620 mol/L, or 0.620 M. This could have also been calculated using the Henderson-Hasselbalch equation, pH = pKa + log([Base]/[Acid]), since the pH, pKa, and [NaOH] were known. The Henderson-Hasselbalch equation also explains why pH = pKa when the NaOH solution neutralized “Ron”. When [Base] = [Acid], [Base]/[Acid] = 1. Taking the log of 1 is equal to 0, leaving the equation to be pH = pKa. A source of error in this experiment could be incorrect calibration of the pH meter. All calculations depended on the recorded pH values given by this tool, so any problems would cause the results to be completely unusable. Conclusion: The purpose of this lab was to use titration curves to graphically determine the molar mass, molar concentration, stoichiometric point, and pKa of the unknown monoprotic weak acid, “Ron”. It was found that the molar concentration of this acid was 0.620M, and the molar mass was 163.7 g/mol. The pKa, which was determined based on the stoichiometric point on the graphs shown in Figures 1-4, was 6.0. The hypothesis was supported since the graphs that illustrated the relationship between volume of NaOH added and the pH of the solution made it possible for all desired values to be calculated. This experiment could have been improved by increasing the amount of trials. This would provide more data to take the average of, which would make the results more reliable. In addition, the trials could have been repeated using multiple different pH meters to dilute the impact of any calibration problems that may have existed. Both of these improvements would ensure more accurate results for this experiment.