Acid Base Titration - CHEM 112 final

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Queens University *

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112

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Chemistry

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Apr 3, 2024

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0 1 Name: Kaliha Dennis Student No: 20420607 Partner: George Ibawli Student No: 20340139 Lab Section:48 Bench # (on computer screen):2 Experiment 10: (2 weeks) (ABT) Acid/base Titration Purpose Part 1 : The Purpose of this experiment is to learn about titration curve, endpoints, and equivalents. In addition, one will learn how to standardize a solution. Part 2: Additionally, the second portion of the experiment to investigate the relationship between the amount of acids and base in a titration and inspect their constants, Understanding the different between the endpoint vs. the equivalence point. Introduction Part I : Monoprotic Acid The process of Titrations is use d to identify the concertation of an unknown solutions. This is usually regarding acids and bases. This is due to the ability to carefully measure the volume of a known solution is slowly added. In this solution there is an indicator that identifies when the titration (or reaction) has gone to compilation. Using the volume of the added solution allows one to determine the unknown solution. The equivalence ne point of a reaction occurs when the number of moles is base added equals the number of moles of acid present. At this point, we can assume that the na, nb, the volumes Va, Vb and the concentrations Ca, Cb the acid and bases following equation. Equation 1: na = VaCa = VbCb= nb The indicator, which identifies the equivalence point is a weak organic acid and has a different color than its conjugate base. The indicator is selected such that at the end point there will be a colour change. It is important not to over titrate the solution, so one can accurately identity the equivalence point This experiment specifically, will determine the concertation of a sodium hydroxide solution using potassium hydrogen phthalate (KHP) , which is dissolved in water. This solution reaction is shown below. Equation 2: KHP -> K+ (aq) + HP – (aq)

0 2 With this, the HP- ions that has formed, when it is mixed with the OH- in the Na)OH solution it will make the solution more basic by getting rid of the OH- Ions, and replacing them with H20, shown in equation 3. This will increase the pH which can be identified by the a light pink pigment. This is due to the chosen indicator, phenolphthalein, which changes colour in the pH range of (8.3-10) . Equation 3: HP- (aq) + OH (aq) - > P2- + H20 (l) Part 2 : Diprotic Acid A Diprotic acid is an acid that produced two H+ ions per molecules when introduced to a solution. Due to this, the acid dissociates in water in two stages. This can be depicted in the two equations bellow. Equation 04: H2X(aq) <-> H+ (aq) + HX (aq) Equation 05: HX- (aq) <-> H+ (aq) + X2- (aq) These two stages are called the two equivalence points. This can additionally be shown using the acid- base reactions occurring between a diprotic acid H2X, and sodium hydroxide base, NaOH are represented below. The start of the reaction H2X (aq) + NaOH (aq) <-> H20(l) + NaHX (aq) The first equivalence point occurs when we take this acid NAHX and use it as its own right with its conjugated base. NaHX (aq) + NaOH (aq) <-> H20 + Na2X (aq) From the start of the reaction we can now produce and identify our second equivalence point. H2X(aq) + 2 NaOH(aq) <-> 2 H20(l) +Na2X(aq) Note that at the first equivalence point, all H- ions from the first dissociation have reacted with base NaOH. At the second equivalence point, all H+ Ions from both reactions have reacted. This in turn, means that the volume of NaOH added at the second equivalence point is twice amount that has been added at the first equivalence. Due to the purpose of the experiment to identify the unknown diprotic acid, one can do this by calculating the molecular weight. With this, one can find the mass of the acid by weighing it. In addition, the concertation can be identified due to equating the number of moles of the unknown acid to the number of moles of the NaOH at the first equivalence point. One the mas and the mole of the diprotic acid is known; the molecular weight can be calculated using the equation below.

0 3 Molecular weight = mass (g) /moles It is additionally possible to determine the acid dissociation constant Ka1 and Ka2 for the two dissociations of the diprotic acid using a half-titration method. Ka1= [H+] [HX-] / [H2x] Ka2 = [H+] [X2-] / [HX-] The first half-titration point occurs when one -half of the H+ ions have been dissociated; therefore [H2X]=[HX-]. Additionally, the second half of the titration point occurs when one-half of the H+ ions in the second dissociation have titrated [HX-]=[X2-]. With this, the pH calculated at the first half titration value is equal the Pka1 value. The first half – titration point volume can be found by dividing the first equivalence point by 2. Likewise, the pH value at the second half titration point is equal to the pHa2 value. The second half titration volume is midway between the first and second. Procedure Note that this procedure has be extracted and reworded for convince from the Queens chemistry First Year Laboratory manual Procedure part I 1. Prepare 50ml burette with ttitrate (0.1M) 2. Pour approximately 10 mL of RO water into the burette, then rinse the burette thoroughly and discard the water down the sink. 3. Repeat the previous step for thorough rinsing And Document the Sample Number of the NaOH Bottle. 4. Fill the burette with around 5-10 mL of NaOH, then rinse the burette and discard the NaOH into your 600 mL waste beaker. 5. Fill the burette with NaOH, ensuring it reaches close to, but not exactly, 0.00 mL (e.g., 1.00 mL is acceptable). 6. Reduce the titrant volume by approximately 0.2-0.5 mL to eliminate any air bubbles at the burette tip. Rinse off any excess NaOH drops from the tip with RO water. 7. Using a clean and dry 250 mL Erlenmeyer flask, weigh approximately 0.8 g of KHP by difference and record the mass to four decimal places. Workload Division Point - Partner 1 (Step 11) and Partner 2 (Steps 12-14) 8. Utilize the molar mass of KHP (204.22 g/mol) to determine the number of moles of KHP in the flask and the estimated volume of NaOH (~0.1 M NaOH). Add roughly 50 mL of RO water to the flask and swirl to dissolve. 9. Introduce 2 drops of phenolphthalein indicator into the flask. 10. Insert a clean stir bar into the flask and place it on the hot plate/magnetic stirrer and Gradually activate the magnetic stirrer to achieve consistent mixing. 11. Position the burette tip slightly below the Erlenmeyer flask neck, ensuring sufficient space to manipulate the stopcock. 12. Add the titrant (NaOH) rapidly until you are approximately 5 mL away from your estimated endpoint.

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0 4 13. Once within 5 mL of the estimated endpoint, slowly add the NaOH in increments of roughly 0.25-0.5 mL. 14. Once 2-2.5ml away add the the NaOH a few drops at a time 15. Once your solution has turned pink 16. Repeat the titration twice more 17. Calculate the actual concentration of NAOh for each trail. Procedure part II Part A 1. Measure out ~0.120 g of the unknown diprotic acid into a 250 mL beaker and record the mass. 2. Add 100 mL of RO water Add a few drops of the chosen indicator to your unknown acid. a. For Bromothymol Blue, use 10-12 drops b. For the other indicators use 5-6 drops 3. Lab Report Note: make note of the color before and after and note also the pH range over which you saw the color change happen Part B: Logger Pro Preparation for pH vs. volume plot (Partner 1) 4. Connect the ph probe via the Go!Link to the USB hub.. 5. Start and set up Logger Pro to plot your pH curve. . a Under Experiment menu > Set Up Sensors > Go!Link: b. Click the square white drop-down box on the right > Choose sensor pH/mV/OR Amplifiers > pH. Exit the window. Change the Data Collection IQ Mode to Events with Entry. c. Configure your point symbols: Part C: Burette Preparation (Partner 2) 6. Prepare the burette by ensuring it is clean. 7. Fill the burette with the approximately 0.1 M NaOH solution. 8. Securely attach the burette to the ring stand using a burette clamp. 9. Empty the NaOH solution into a 400 mL waste beaker until it reaches the 0.00 mL mark. 10. Precisely record the initial volume reading to two decimal places from the burette. 11. Position a paper towel on the stir plate beneath the beaker to facilitate endpoint detection. 12. Position the beaker on the magnetic stirrer/hot plate and gradually activate the magnetic stirrer to achieve a consistent mixture. 13. Put pH probe into the acid solution using a utility clamp on a ring stand as depicted in the provided figure. 14. Place the probe aside from the beaker to prevent any disruption from the stir bar. 15. Position the burette so that its tip is submerged inside the beaker with sufficient space to manipulate the stopcock. Part D: Titration Obtain data on the pH vs. Volume curve during the titration process. Take note of the endpoint, marked by the volume and pH where the indicator's color changes. 16. Start data collection by clicking on "Collect" in Logger Pro 17. Begin adding the titrant gradually, approximately 1 mL at a time. After each addition, click "Keep" and input the volume added up to that point in the titration. A data point will be generated on the graph. 18. Upon completing the titration, click on "Stop" to halt data collection.

0 5 19. Rescale the graph by selecting "Autoscale LA" or manually adjusting the axes. Part E: Plotting the Second Derivative on Logger Pro 20. Note the most distinctly defined equivalence point for later reference. If two equivalence points are closely situated, it's likely due to similar pKa values (<2 units apart). 21. Plot the second derivative curve and determine the point where it intersects zero a. In Logger Pro, navigate to Page > Add Page. Choose "Copy Current Page" from the dialogue box b. Go to the Data menu and select "New Calculated Column." c. Title the graph "Second Derivative" in the text entry box. d. For the equation, select Functions > Calculus > Second Derivative. Then select Variables (Columns) > pH. 22. Take a picture of your graphs Data and Observations Part I: Monoprotic Acid Data Table Trail 1 Trial 2 Trail 3 Mass of Weighing containor (empty) 106.9407 g 114.283 g 114.252 g Mass of container and mass of KHP 107.7475g 113.4278g 113.4513g Mass of KHP 0.8068 g 0.8552 g 0.8007 g Moles of KHP 3.951 x 10 ^ -3 mol 4.1876 x 10 – 3 mol 3.92x10^-3 mol Initial Burette volume (mL) 1.2 ml 1.3 ml 1.7ml Final Burette volume (mL) 41.2ml 43.6ml 41.2ml Total Volume of NaOH ( Vfinal- Vintital) ml 40ml 42.3 ml 39.5 ml Concentration of NaOH 0.098775 0.098997 0.100255 Observations

0 6 - In trail 1 the colour changed a slight pink colour at a titration of NaOH of 39.5ml - In trail 2 the colour additionally changed a slight pink colour which was difficult to see until there was a white surface behind the solution. This slight changed occurs at a volume of 41.9ml - Lastly, there was a change at 39.2ml. Part II: Monoprotic Acid Sample Number: 2 pH vs Volume of NaOH graph Second Derivative Graph Observations made during the laboratory - The colour change occurred at 28.6ml with a ph of 7.1

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0 7 Questions Questions for Part I 1. What is the concentration of the base from the stock bottle? Show any calculations and uncertainty calculations. Trial 1: Mass of KHP in flask = 0.8068 g Volume of NaOH added to flask = 41.2ml Calculation of molarity of NaOH for trial 1: Molar mass of KHP =204.23 g/mol 0.8068 g/204.23 g/mol=0.00395 mol of KHP The mol of KHP will equal to 0.00395 NaOH Therefore the Molarity of NaOH for trail 1=0.00395mol/0.0412 L= 0.098775 M Trial 2: Mass of KHP in flask = 0.8552g Volume of NaOH added to flask = 43.6ml Calculation of molarity of NaOH for trial 1: Molar mass of KHP =204.23 g/mol 0.8552 g/204.23 g/mol=0.004187 mol of KHP The mol of KHP will equal 0.004187 NaOH Therefore, the Molarity of NaOH=0.00418mol/0.00436 L= 0.09897 M Trial 3: Mass of KHP in flask = 0.8007 g Volume of NaOH added to flask = 41.2ml Calculation of molarity of NaOH for trial 1: Molar mass of KHP =204.23 g/mol 0.8007 g/204.23 g/mol=0.00382 mol of KHP The mol of KHP will equal 0.00382 NaOH Therefore, the Molarity of NaOH=0.00382mol/0.00412L= 0.1002M Average concertation = Concentration trail 1 + Concentration Trail 2 + Concentration Trial 3 = (0.098775 + 0.098997 + 0.100255)/ 3 = 0.09934

0 8 Therefore, the average concertation of NAOH Is 0.09 mol/L Questions for Part II 2. Calculate the molecular weight of the diprotic acid in g/mol using the mass of diprotic acid that you measured in the first step of the procedure and the moles you determined from the titration results (Show all your work). Finding the Molar Mass of the Unknown Acid Mass of Diprotic Acid= 0.172g The concentration of NaOH used = 0.100M The volume required to neutralize the acid = 13.45ml Finding the Moles of the NaOH Solution (0.100 x (13.45 x 10 ^-3) = 1.345 x 10 ^-3 Moles of NaOH = 1.345 x 10 ^-3 Finding the Moles of Unknown Acid Moles of NaOH = = 1.345 x 10 ^-3 = Moles of unknown acid Therefore, the Moles of the Unknown acid is = 1.345 x 10 ^-3 +/-0.001 mol Finding the Molar Mass of Unknown Acid Molar mass = Mass / moles = 0.172/ (1.345 x 10 ^-3) = 127.88 g/mol 3. From the following list of five diprotic acids, identify the unknown diprotic acid that you used. Note that not everyone uses the same acid. Diprotic Acid Formula MW Oxalic Acid H 2 C 2 O 4 ·H 2 O 126 Malonic Acid H 2 C 3 H 2 O 4 104 Maleic Acid H 2 C 4 H 2 O 4 116 Malic Acid H 2 C 4 H 4 O 5 134 Tartaric Acid H 2 C 4 H 4 O 6 150 4. Determine the percent error between the molecular weight value you calculated in step 2 and that of your chosen acid. = [(127.88-126)/126] x 100 = 1.49 % error 5. Determine the pH for the first half-titration point (½ of the first equivalence point volume) and the second half-titration point (halfway between the first and second equivalence points). Compare these values to the pK a1 and pK a2 values for your chosen acid (reference). First half titration point - Ph1 = 1.9 - Pka1 = 1.27 1

0 9 - Ph2 = 3.8 - Pka2 = 4.27 1 6. By comparing various indicator results with other groups around you, which indicator(s) do you think worked for this experiment, i.e., the colour changed very near the equivalence point. I believe that the best indictor with the groups around me in this case would be methyl red. The reason being is because methyl red represents its colour when the pH of solution around 2.8 -3.8. With this, this closely falls in the range of the two titration points of the diprotic acid. If one was given the variety to choose an indicator, that was not limited to the options given, I would choose Congo red because Congo red has an indication range that is greater, allowing for it to include the pH of the second titration point. Thus, it has a pH range of 2.0- 5.0 2

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0 10 References 1. National Center for Biotechnology Information (2024). PubChem Compound Summary for CID 971, Oxalic Acid. Retrieved February 18, 2024 from https://pubchem.ncbi.nlm.nih.gov/compound/Oxalic-Acid. 2. Queen's University, & Department of Chemistry. (2023). First-Year Laboratory Manual (2023- 2024 ed.).