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Acid Base Titration Background Acid-Base titrations are examples of volumetric analysis where the concentration of an un- known acid or base can be determined precisely from a known amount of the reactant (either acid or base). In an acid-base titration, one begins with the stoichiometrically balanced chemical equation which is assumed to go to completion. The equivalence point, then is the point during a titration when acid (or base) is completely neutralized by adding a stoichiometric amount of base (or acid). In other words, at the equivalence point the moles of H+ are equal to the number of moles of OH−. The equivalence point is experimentaly determined by using an indicator. The point at which the color of the indicator changes from one color to another color and is known as the endpoint of the titration. Before we do any acid-base titration, we must accurately determine the concentration of either acid or base with which we can carry out the subsequent titration. This process is called the ”Standardization” of the acid or base. In this experiment, we will standardize a solution of NaOH. Part A: Standardization of a NaOH Solution. We will use a primary standard such as ”KHP” (potassium hydrogen phthalate) to determine the exact concentration of NaOH solution according to the following chemical equation: NaOH(aq) + KHC8H4O4(aq) −→ NaKC8H4O4(aq) + H2O(l) At the equivalence point, the number of moles of NaOH is equal to the number of moles of KHP because they react in a 1:1 mole ratio. For each of the three trials use the mass of KHP and the volume of base, NaOH solution, needed to reach the endpoint to calculate the molar concentration of the NaOH solution. If the calculated NaOH concentrations do not agree within 0.8% do a fourth trial and average the results. 1 Part B: Percent by mass of Acetic Acid in Commercial Vinegar The percent by mass of acetic acid, CH3CO2H, in vinegar is given by the equation shown below. Percent by mass CH3CO2H = mass of CH3CO2H mass of solution ∗ 100 The mass of acetic acid (g) can be calculated if you know the molarity (M) of the acetic acid in vinegar, the density of vinegar (densityvinegar=1.005 g/mL), and the molar mass of the acetic acid (HC2H3O2, 60.06 g/mol). The molarity of acetic, CH3CO2H, is determined by titrating a known volume of CH3CO2H with the NaOH solution standardized in Part A of the experiment. The reaction between CH3CO2H and the NaOH solution is given by the equation below. CH3CO2H(aq) + NaOH(aq) −→ NaCH3CO2(aq) + H2O(l) 2 Pre-Lab Name: Lab Section 1. The reaction between KHP (potassium hydrogen phthalate, KHC8H4O4) and NaOH is given by the following chemical equation. Write your answer in the space provide below. NaOH(aq) + KHC8H4O4(aq) −→ NaKC8H4O4(aq) + H2O(l) Use the experimental data in the table below to calculate the Molarity of the NaOH solution. Mass of KHP Used 1.198g Initial Volume of NaOH 0.45 mL Final Volume of NaOH 19.75 mL Molarity of NaOH solution = 2. The reaction between nitric acid, HNO3(aq), and NaOH is shown below. HNO3(aq)(aq) + NaOH(aq) −→ NaNO3(aq) + H2O(l) Use the experimental data provided below and the molarity of the NaOH solution determined in question 1 to calculate the Molarity of the HNO3 solution. Volume of HNO3 sample 10.00 mL Initial Volume of NaOH 1.58 mL Final Volume of NaOH 24.73 mL 3 Experimental Procedure Part A: Standardization of a NaOH Solution 1. Start Virtual ChemLab, select Acid-Base Chemistry, and then select Acid-Base Stan- dardization from the list of assignments. The lab will open in the Titrations laboratory. 2. Click the Lab Book to open it. Click the Beakers drawer and place a beaker in the spotlight next to the balance. Click on the Balance area to zoom in, open the bottle of KHP by clicking on the lid (Remove the Lid). Drag the beaker to the balance to place it on the balance pan and tare the balance. Pick up the Scoop and scoop out some sample by first dragging the scoop to the mouth of the bottle and then pulling the scoop down the face of the bottle. As the scoop is dragged down the face of the bottle it will pickup different quantities of solid. Select the largest sample possible and drag the scoop to the beaker on the balance until it snaps in place and then let go. Repeat this one additional time so you have put two scoops (approximately 2 g) of KHP in the beaker. Record the mass of the sample in the data table on the following page and return to the laboratory. 3. Drag the beaker from the balance to the sink and hold it under the tap to add a small amount of water. Place it on the stir plate and drag the calibrated pH meter probe to the beaker. Add Phenolphthalein as the indicator. 4. The buret will be filled with NaOH. Click the Save button in the Buret Zoom View window so the titration data can be saved. The horizontal position of the orange handle is off for the stopcock. Open the stopcock by pulling down on the orange handle. The vertical position delivers solution the fastest with three intermediate rates in between. Turn the stopcock to one of the fastest positions. Observe the titration curve. When the blue line begins to turn up, double-click the stopcock to turn it off. Move the stopcock down one position to add NaOH solution drop by drop. 5. Repeat at two additional times. For each trial be sure to record the mass of KHP and the volume of NaOH needed to reach the endpoint. Do not forget to refill the buret with NaOH solution and place the pH meter and indicator in the beaker each time. The molar mass of KHP is 204.22 g/mol. 6. Calculate the molarity of NaOH to four decimal places. Three trials should agree within ±0.8% of the mean value. If not, do a fourth trial. (You may discard one outlying Trial if you are certain it is not within ±0.8% of the mean). In your the results section of your final report, you will include the average Molarity and average deviation from the mean. 7. When you are done. Click on ”Clean up Lab Bench.” This is the red bucket on the right side of the lab bench. 4 Part B: Percent by mass of Acetic Acid 1. Go to the stockroom and click on the ”Unknowns” label on the bottom shelf. 2. For the unknown bottle on the left side. Click on the arrow just to the right of ”Select Unknown.” Scroll down the list of the unknown and select Acetic Acid. 3. Now set the concentration range for the Acetic Acid solution. Enter 0.9000 for the MAX Concentration. Then enter 0.7000 for the MIN Concentration. 4. For the unknown bottle on the right side. Select Sodium Hydroxide. In the MAX Concentration enter the average NaOH Concentration you determined in Part A on this experiment. This solution, of NaOH, is now your standard solution. 5. Click on the Buret to open the Buret Zoom view. Fill the buret with your standard NaOH solution. 6. Use the 10.00 mL pipet to deliver a sample of the Acetic Acid solution into a beaker. 7. Turn On the pH meter and calibrate it with the pH 4 and pH 10 solutions. 8. On the Buret Zoom View click Graph. This will create a graph of the titration curve. The click Save on the Buret Zoom View. This will save your titration data. 9. The horizontal position of the orange handle is off for the stopcock. Open the stopcock by pulling down on the orange handle. The vertical position delivers solution the fastest with three intermediate rates in between. Turn the stopcock to one of the fastest positions. Observe the titration curve. When the blue line begins to turn up, double-click the stopcock to turn it off. Move the stopcock down one position to add NaOH solution drop by drop. 10. Perform two additional trials. 11. Calculate the molarity (M) of Acetic Acid to four decimal places. Your trials should agree within 0.8% of the mean value. If not complete one more trial. You may discard one outlying trial if you are certain it is not within the mean value. In the Results section of your report, you will be asked to report the average Molarity and average deviation from the mean. 12. Using the average molar concentration of Acetic Acid, calculate the percent by mass of Acetic Acid. The density of the Acetic Acid solution is 1.005 g/mL. 5 Data Sheet Part A. Standardization of NaOH solution Trial 1 Trial 2 Trial 3 (a) Mass of KHP used (g) (b) Initial volume of NaOH (mL) (c) Final volume of NaOH (mL) (d) Volume of NaOH used (mL) (e) Molarity of NaOH solution (f) Average Molarity of NaOH solution = (g) Average deviation from the mean value ( δ ) = (h) Percent average deviation from the mean value ( δ x ) ∗ 100 = Part B. Percent Acetic Acid in Vinegar Unknown ID: Trial 1 Trial 2 Trial 3 (a) Volume of vinegar used (mL) (b) Initial volume of NaOH (mL) (c) Final volume of NaOH (mL) (d) Volume of NaOH used (mL) (e) Molarity of CH3CO2H (f) Average Molarity of CH3CO2H = (g) Average deviation from the mean value ( δ ) = (h) Percent by mass acetic acid in vinegar = 6 Post-Lab Questions 1. If after weighing the KHP sample some of the solid KHP powder was spilled while transferring the beaker from the balance to the Stir Plate. Would the calculated molarity of the NaOH solution be: Too High, Too Low, or Unaffected? Explain. 2. If liquid splashed out of the Erlenmeyer flask during the titration but before the end- point was reached, would the calculated molarity of the NaOH solution be: Too High, Too Low, or Unaffected?. Explain. 3. If air bubbles were not expelled from the base buret tip and remained in the buret prior to titration, would the calculated molarity of CH3CO2H in Acetic Acid be: Too High, Too Low, or Unaffected?. Explain. 7