Microsoft Word - Laboratory 3 Titration 1 Name ________________________________ Date _Fall 2021_____ Laboratory 3, Calculation of Molarity of H3PO4 by Titration with NaOH Chemistry 201, Wright...

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Microsoft Word - Laboratory 3 Titration 1 Name ________________________________ Date _Fall 2021_____ Laboratory 3, Calculation of Molarity of H3PO4 by Titration with NaOH Chemistry 201, Wright College, Department of Physical Science and Engineering Molarity is a common unit within the chemical laboratory to denote concentration of a solution. The units of molarity are moles of solute per liter of solution, mol/L, or commonly, M. In this experiment, you will determine the concentration (molarity) of an unknown phosphoric acid solution, H3PO4. The concentration of the phosphoric acid solution will be determined via titration using NaOH with a known molarity. The balanced chemical equation will follow, H3PO4 (aq) + NaOH (aq) ---------> NaH2PO4 (aq) + H2O (l) At the end point, the moles of H3PO4 will equal the moles of NaOH. You will use Bromocresol Green as the indicator. Bromocresol Green is yellow colored in acidic solution and is blue in basic solution. The end point of the reaction will be a green solution. Watch video : Titration with bromocresol green https://www.youtube.com/watch?v=OnHMp9Rn2gI PRELAB ASSIGNMENT 1. What are the units of molarity? 2. What is the acid used in today’s titration? 3. What is the base used in today’s titration? 4. What is the purpose of today’s experiment? 5. Which indicator will be used in this experiment? 6. What color is the indicator in H3PO4 solution? 7. What color will the indicator be at the end point of the reaction? 2 PROCEDURE All waste can be put down the sink. 1. Obtain a burette which will hold the sodium hydroxide (NaOH) solution. Be very careful with the burette, it is very expensive. 3. To prepare the burette for titration: obtain about 200 mL of NaOH solution. Use about 25 mL of the NaOH solution to rinse out the NaOH burette. Don’t forget to release some of the NaOH solution out the bottom of the burette. This is waste, put it down the sink. Record the molarity of the NaOH solution (1). 4. Fill the burette close to the zero line (but not above) with the NaOH solution and record the initial volume (3). Remember to read the burette to 0.01 mL. 5. Obtain 22.0 mL of H3PO4 solution from the dispenser into a clean 250 mL Erlenmeyer flask. 6. Add about 25 mL of deionized water and 3 drops of Bromocresol Green indicator to the Erlenmeyer flask containing the H3PO4. Swirl the Erlenmeyer flask to make sure all components are mixed. 7. Place a piece of white paper under the NaOH burette. This will be used to observe colors and color changes. Bromocresol Green indicator appears yellow in acidic solution and appears blue in basic solution. The end point for this titration is a green color (right between yellow and blue). 8. Start adding NaOH to the Erlenmeyer flask containing the H3PO4 solution. The solution starts out as yellow (due to the phosphoric acid), but you should start to see a blue color as the base (NaOH) is added. Swirl the Erlenmeyer flask as you are adding the NaOH solution. Start slowing down the addition of the NaOH, as you don’t want to overshoot the green endpoint. Add the NaOH in small increments, or even, dropwise, when you are close to the endpoint. Stop adding NaOH when the color of the solution is green. 9. When you have reached the endpoint, read the burette to the 0.01 mL and record the final volume of NaOH (4). Calculate the volume of NaOH in the Erlenmeyer flask by subtracting the final minus initial NaOH volumes (5). 10. The amount of H3PO4 in the solution is related to the amount of NaOH necessary to reach the endpoint. Convert the mL of NaOH used to L of NaOH (6). Show your calculation for Trial 1 here. 3 11. Use the molarity of the NaOH to convert L of NaOH to moles of NaOH (7). Show your calculation for Trial 1 here. 12. Calculate the moles of H3PO4 that reacted (8). Consult the coefficients in the balance chemical reaction to obtain the mole ratio. Show your calculation for Trial 1 here. 13. Convert the mL of H3PO4 to L of H3PO4 (9). Show your calculation for Trial 1 here. 14. Finally, calculate the molarity of the H3PO4 solution by dividing the moles of H3PO4 by the liters of H3PO4 (10). Show your calculation for Trial 1 here. 15. Refill the burettes with their solutions and repeat the titration 2 more times. Calculate and record the average molarity (11). 4 Calculation of molarity of H3PO4 Report Sheet 1. Molarity of the NaOH solution _0.238__ mol/L Trial 1 Trial 2 Trial 3 2. Volume of H3PO4 added to flask 22.0 mL 22.0 mL 22.0 mL 3. Initial NaOH volume _0.25___ mL __0.75_ mL _0.55___ mL 4. Final NaOH volume _17.60__ mL __18.20_ mL __17.90_ mL 5. NaOH volume used for titration to reach green end point ________ mL ________ mL ________ mL 6. NaOH volume used for titration ________ L ________ L ________ L 7. Moles of NaOH used for titration ________ mol ________ mol ________ mol 8. Moles of H3PO4 that reacted ________ mol ________ mol ________ mol 9. Volume of H3PO4 added to flask ________ L ________ L ________ L 10. Molarity of H3PO4 ________ mol/L ________ mol/L ______ mol/L 11. Average molarity of H3PO4 ________ mol/L 5 QUESTIONS (Show your work) 1. Calculate the molarity of the solution prepared from 0.45 g Li2SO4 in 250. mL of solution. 2. According to the reaction below, how many grams of precipitate will be produced from 12.39 mL of 0.258 M CaCl2 solution? Assume excess Na3PO4. 3 CaCl2 (aq) + 2 Na3PO4 (aq) --------> Ca3(PO4)2 (s) + 6 NaCl (aq) 6 3. According to the equation below, how many milliliters of 0.32 M Na3PO4 will react with 11.25 mL of 0.46 M AgNO3? Na3PO4 (aq) + 3 AgNO3 (aq) --------> Ag3PO4 (s) + 3 NaNO3 (aq) 4. Using the following reaction, H2SO4 (aq) + 2 NaOH (aq)  Na2SO4 (aq) + 2 H2O (l) calculate the molarity of the H2SO4 solution if 14.92 mL of NaOH was necessary to reach the endpoint of a titration. The molarity of the NaOH solution was 0.83 M and 25.18 mL of H2SO4 was added to the Erlenmeyer flask. 7
Oct 19, 2021
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