Experiment 3 acid-base titration-Report PDF

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CHEM 1202 Expt. 3 Acid-Base Titrations Waniya Sheikh

Introduction: Visual endpoint titration is a method of quantitatively analyzing the endpoint of an acid-base neutralization. At the endpoint, no more acid can react with the base (or vice versa) in the neutralization. By using an indicator, the individual performing the titration will be able to detect the endpoint via a change in solution color as the solution is at its endpoint. In the case of this experiment, phenolphthalein is used, so once the neutralization reaches its endpoint, the solution will turn pink. Phenolphthalein is a weak, colorless acid, while its ion is a bright pink. As more hydroxide ions are added through a titration, the hydroxide reacts with the hydrogen ions, therefore neutralizing the acid. Once there is an excess of hydroxide ions in the solution, the pH of the solution increases, thus the equilibrium of the solution shifts to form more phenolphthalein ions. Potentiometric acid-base titration uses a pH meter to monitor the pH in the solution being titrated. If there is a change in concentration of an ion, the electrode of the pH monitor would detect the change in potential in the solution. This method of titration is often used when there is no suitable indicator for the titration, and is preferred in titrations involving polyprotic acids. By relying on the pH values recorded rather than the color of an indicator to determine the endpoint of a reaction, potentiometric titrations are generally more precise in determining the endpoint in a solution 2. In using these pH values, a titration curve is then plotted, which is used to determine the equivalence point of the reaction. In this experiment, titrations of a strong acid and strong base, as well as of a weak acid and strong base are carried out. The strong acid and strong base neutralization is represented in this equation: HCl(aq)+NaOH(aq)  NaCl(aq)+H2O(l) The reaction between the weak acid and strong base is represented in this equation: KHC8H4O4(aq)+NaOH(aq) KNaC8H4O4(aq)+H2O(l) The objective of this experiment is to determine the visual endpoint of a strong acid vs. strong base neutralization, to determine the potentiometric endpoint of a strong acid vs strong base neutralization, and to determine the Ka of KHP via potentiometric titration. Experimental: Part A of the experiment dealt with the determination of a visual endpoint between a strong acid and strong base using an indicator. First, 10 mL of HCl was pipetted into a 250 mL Erlenmeyer flask, and then diluted with approximately 50 mL of water. Three drops of phenolphthalein were then added to the diluted solution. A burette was rinsed and then filled with NaOH solution and the initial burette reading †

I would like to acknowledge my lab partner _________________ for their assistance with the experiment.

was recorded. As the titration began, NaOH was slowly added from the burette into the flask while swirling the flask. The NaOH was then added drop-wise until the solution in the flask turned pink. Once the solution turned pink, the final burette reading was recorded. This titration was repeated two more times. Part B dealt with the determination of a potentiometric endpoint between a strong acid and a strong base using a pH meter. 10 mL of HCl was first pipetted into a 150 mL beaker, and then diluted with approximately 50 mL of water. A magnetic stirrer was then dropped into the beaker. The pH meter was then calibrated and the electrode was fully immersed into the solution. Using the same NaOH from part A, the burette was refilled, and the initial volume of NaOH was recorded. Before the titration, the pH of the initial solution was recorded. The titration was started by adding 2 mL of NaOH into the solution, recording the exact volume of NaOH and pH after each addition until 5 mL short of equivalence point (which was determined visually in part A). 1 mL additions of NaOH were then added until approximately 2mL short of equivalence point. At this point, NaOH was added dropwise until the approximately 2 mL past equivalence point. The titration was continued in 1 mL and 2 mL increments for approximately 10 more Mr. The pH and volume of NaOH was recorded after each addition of NaOH. This whole process was performed twice. Part C dealt with the determination of the Ka of KHP through a potentiometric titration a weak acid and a strong base. 3.9856 g of KHP was weighed out and put into a 150 mL beaker, which was then dissolved with approximately 50 mL of water. Following the same procedure as in part B, after each addition of NaOH, the pH and the volume of NaOH was recorded. The KHP solution was slowly titrated to approximately half-equivalence point and then to its equivalence point. Additional NaOH was titrated into the KHP solution approximately 8 mL past equivalence point. This titration was conducted twice. Results: Concentration of NaOH: _________________________________ Part A: Strong acid vs strong base – Visual end point using indicator Burette readings Burette reading( NaOH) Final (mL) Initial (mL) Volume used (mL) Average volume(NaOH) used (mL)

Trial 1 21.50 0.25 21.25

Trial 2 39.90 21.25 18.65 19.67

Trial 3 37.50 18.40 19.10

Table 1: Burette readings for strong acid vs. strong base visual endpoint determination using indicator.

Part B: Strong acid/ strong base– Potentiometric end point using pH Trial 1

Trial 2 2

Burette reading (mL) 0 2.00 4.00 6.00 8.00 10.00 12.00 13.00 14.00 15.00 16.00 16.50 17.00 17.50 18.00 18.50 19.00 20.00 21.00 22.00 23.00 24.00

NaOH added(mL) 0 2.00 4.00 6.00 8.00 10.00 12.00 13.00 14.00 15.00 16.00 16.50 17.00 17.50 18.00 18.50 19.00 20.00 21.00 22.00 23.00 24.00

Measured pH 1.05 1.21 1.16 1.10 1.12 1.14 1.19 1.20 1.25 1.32 1.44 1.52 1.61 1.78 2.21 10.36 11.49 11.97 12.18 12.30 12.39 12.45

Burette reading (mL) 0 2.00 4.00 6.00 8.00 10.00 12.00 13.00 14.00 15.00 16.00 16.50 17.00 17.50 18.00 18.50 19.00 20.00 21.00 22.00 23.00 24.00

NaOH added (mL) 0 2.00 4.00 6.00 8.00 10.00 12.00 13.00 14.00 15.00 16.00 16.50 17.00 17.50 18.00 18.50 19.00 20.00 21.00 22.00 23.00 24.00

Measured pH 1.00 0.95 0.96 0.98 1.03 1.09 1.17 1.21 1.33 6.73 11.46 11.43 11.61 11.66 11.74 11.78 11.90 12.03 12.19 12.26 12.33 12.39

Table 2: Burette readings for strong acid vs. strong base measuring end point using pH.

Part C: Weak acid/ strong base-Potentiometric endpoint using pH. Trial 1 4.5696 0.6112 3.9584

Mass (weigh dish+KHP) (g) Mass (weigh dish+residue) (g) Mass (KHP) (g)

Trial 2 4.6120 0.6315 3.9805

Table 3: Recorded masses of KHP in dish, KHP residue in dish, and total mass KHP used.

Weak acid/ strong base – Potentiometric end point using pH

Burette reading (mL) 0 2.00 4.00 6.00 8.00 10.00 11.00 12.00

Trial 1 NaOH added(mL) 0 2.00 4.00 6.00 8.00 10.00 11.00 12.00

Measured pH 4.44 4.28 4.49 4.70 4.80 5.02 5.17 5.26

Burette reading (mL) 0 2.00 4.00 6.00 8.00 10.00 11.00 12.00

Trial 2 NaOH added (mL) 0 2.00 4.00 6.00 8.00 10.00 11.00 12.00

Measured pH 4.43 4.62 4.83 5.02 5.16 5.35 5.44 5.53 3

13.00 13.50 14.00 14.50 15.00 15.50 16.00 16.50 17.00 17.50 18.00 18.50 19.00 19.50 20.00 21.00 22.00 23.00 24.00 25.00

13.00 13.50 14.00 14.50 15.00 15.50 16.00 16.50 17.00 17.50 18.00 18.50 19.00 19.50 20.00 21.00 22.00 23.00 24.00 25.00

5.34 5.40 5.58 5.63 5.69 5.82 5.91 5.95 6.20 6.21 6.59 9.85 9.18 10.30 11.78 12.60 12.72 12.79 12.88 12.90

13.00 13.50 14.00 14.50 15.00 15.50 16.00 16.50 17.00 18.00 19.00 20.00 21.00 22.00 23.00 24.00

13.00 13.50 14.00 14.50 15.00 15.50 16.00 16.50 17.00 18.00 19.00 20.00 21.00 22.00 23.00 24.00

5.67 5.73 5.79 5.86 5.93 6.03 6.15 6.48 6.88 10.44 11.83 12.14 12.31 12.44 12.53 12.54

Table 4: Burette readings for weak acid vs. strong base measuring end point using pH.

Discussion: The objectives of this experiment were met by successfully determining the visual endpoint of a HCl and NaOH both visually and by means of potentiometric titration, as well as determined the Ka of KHP via potentiometric titration. Based on the results in part A of the experiment, the visual endpoint of an HClNaOH neutralization came to be around 19.67mL of NaOH used. Comparing this value to the endpoint volume in part B, the volumes are very close to each other, as the spike in measured pH was around 18.5 mL. The phenolphthalein used in part A to indicate the change in pH was less precise than measuring via pH meter, but was appropriate given the range of the indicator (phenolphthalein has a large pH range from 8.3 to 10.0). Such a wide pH range makes phenolphthalein a popular indicator in acid-base titrations. However, the drawback to using an indicator with such a wide range is less precision; albeit not a large difference, the equivalence point volume differed in 0.24 mL. The results from part C was of a titration between a weak acid and a strong base. Compared to the results in part A and B, there is a difference in how the reaction takes place in part C. Since a weak acid is being used, the pH of the weak acid will be much higher than in part A and B. Furthermore, there is not a huge jump in pH change when equivalence point is reached, but rather just a difference in pH of 4.90 between 17.5 mL and 18.5 mL, compared to the difference in pH of 8.24 between 18.0 mL and 18.5 mL. This is further evident when the titration curves were graphed. In comparison to part C, the curve for part B was not only larger (starting off at a much lower pH), but the slope at the equivalence point was much steeper compared to a modest steepness at equivalence in part C. The reason why the equivalence point in part C occurs at a pH higher than 7.00 (in this case, it was found to be at 9.00) is because this is a titration between a weak acid and a strong base. Because weak acids have conjugate 4

bases, there is a lag (or a buffer zone before the equivalence point) in which some of the weak acid converts into its conjugate base. This results in a larger amount of NaOH to be used in order for the reaction to reach its equivalence point. A major source of error that contributed to the large percent error when attempting to determine the experimental molar mass of KHP came from the incomplete dissolving of the KHP crystals in the water. After waiting several minutes for the crystals to dissolve, part C of the experiment was carried, despite some of the crystals sitting at the bottom of the beaker. This resulted in a lower than expected concentration of KHP solution when carrying out the titration, thus giving a much lower experimental molar mass (compared to the theoretical value of 204.22 g/mol) and a high percent error when calculating. In further experiments, it is to be made sure that sufficient time is given to dissolve as much of the KHP crystals as possible before carrying out the titration.

References: 1

Clark, Jim. Acid-Base Indicators. http://www.chemguide.co.uk/physical/acidbaseeqia/indicators.html (accessed Oct 13, 2014) 2 Mettler-Toledo International Inc. FAQ Titration (II) http://us.mt.com/us/en/home/supportive_content/product_information_faq/FAQ_Titration. ml (accessed Oct 13, 2014) 3 Department of Chemistry. CHEM 1202 Lab Manual; Mount Royal University: Calgary, Fall 2014.

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