Q EXPERIMENT 17: ACID-BASE REACTIONS AND TITRATION Introduction: This experiment demonstrates the analytical technique, titration. In a titration, a solution is delivered from a burette until it completely consumes another solution in a flask. You will chemically react an acid with a base, until the endpoint of the reaction is shown by the indicator phenolphthalein. You will first determine the concentration of a base solution (standardization) and then of an unknown acid solution using previously standardized base. The reaction of a strong acid and a strong base goes quickly: i.e. HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l) acid base salt water This reaction is often called neutralization. Part A: To standardize the NaOH solution, we will react it with potassium hydrogen phthalate, KHC8H4O4 Because of its complex formula, this compound is often called “KHP”. The molar mass of KHP is 204.2 grams per mole. KHC8H4O4(aq) + NaOH(aq) KNaC8H4O4(aq) + H2O(l) Starting with a known mass of KHP then recording the volume of NaOH needed to reach the endpoint, we can calculate the molarity of the base .We want the titrated solution to be a very pale pink, not bright rosy red, at the endpoint. Part B: Once the concentration of the NaOH solution in the buret is known, we can determine the concentration of an HCl solution by titrating it with the NaOH solution from part A. Phenolphthalein is used as the indicator by adding it to the HCl solution. Materials needed: Chemicals Equipment Solid KHC8H4O4 (“KHP”) Deionized water Phenolphthalein solution Approximately 0.2M NaOH solution Unknown HCl solution Weighing vial , scoop or spatula 3 125-mL Erlenmeyer flasks or 2 flasks and a 250 mL beaker wash bottle with deionized water. 50-mL buret, buret holder and ring stand Plastic funnel. 500-mL florence flask Procedure: Part A: Standardization of an unknown base 1. Obtain and rinse a burette with deionized water from your wash bottle. CAUTION: do not stick the end of the burette under the faucet! It leaks all over the floor. Be sure to open the valve at the bottom and let some water rinse out the stopcock. If you think the tip is clogged, ask your instructor to check it. 2. Practice reading the meniscus in your burette while cleaning it with deionized water. Read your burette by estimating between the 0.1-mL marks. In other words, your recorded volume measurements include an uncertain digit at 0.01 mL. If the meniscus is right on a mark, record the second decimal place as a zero. Hold the clean burette in the burette clamp that is attached to the ring stand. 3. Obtain a 500 mL Florence flask, clean it well with deionized water, and then pour approximately 16-17 mL of 6 M NaOH. Ask your instructor which bottle of sodium hydroxide your need to use. Dilute this solution with deionized water to just below the neck of the flask. Do not worry about exact measurements; you will use the technique of standardization to determine the exact concentration. Mix the solution well. Do not discard this solution until the entire experiment is complete. 4. When your burette is clean, rinse it 2 times with small portions (about 15- 20 mL) of your NaOH, discarding the rinses. Then clamp the burette in on to a burette clamp. Close the valve at the bottom, place a plastic funnel in the top opening and carefully pour base solution into the burette until the solution level is near the 0.00 mL mark). Make sure there are no air bubbles trapped in the tip of the burette. Record the initial base volume reading for this trial. (0.00 or slightly below is OK) 5. Weigh three samples of KHP (about 1.00 to 1.10g into three separate 125-mL Erlenmeyer flasks and/or beakers. Label the flasks and record your weights. 6. Dissolve the KHP crystals in your flasks in about 30 mL of deionized water if some KHP sticks to the sidewalls of the flask; wash it down with deionized water from your wash bottle. If the KHP doesn’t dissolve in a short time, you may gently warm the solution in a hot tap water bath. Add 2-3 drops of phenolphthalein solution from the dropper bottle on the shelf. 7. Place the flask under the burette. A piece of white paper under the flask makes it easier to see the pale pink color at the endpoint. Open the valve and allow base to flow from the burette into the flask. Swirl continually to mix the solutions. As you get close to the endpoint, the solution will begin to show pink color that goes away when you mix. Slow the rate of base addition to one drop at a time, If you splash the solution up onto the sidewalls of the flask, spray a stream of water from your wash bottle over the inside of the flask. The extra water will not affect your results. When one drop of base changes the solution from colorless to pale pink, close the burette valve, rinse down the flask one last time, and make sure that the pink color lasts for at least 30 seconds. If so, record the final burette volume reading. 8. Discard the titrated solution into the sink. Do at least three successful titrations that achieve a pale pink color of the indicator. If color is bright rosy red, you have overshot the endpoint and cannot include that trial in your calculations. If you have time after completing the standardization of the base, you may continue right into Part B, or keep the NaOH solution, tightly stoppered, in your locker or in the cabinets in the back of the room, until the next lab. 9. At the end of the lab period, drain and discard the leftover base solution from the burette and rinse the burette with several portions of water. Return the burette. Part B: Determination of an unkown acid 1. Prepare a burette, filled with standardized NaOH solution, and three clean 125-mL Erlenmeyer flasks as you did earlier. 2. Your instructor will assign you one of the several unknown HCl solutions for your experiment. Be sure to record the unknown code in your notebook and on your lab report. Pour about 100 mL of your assigned unknown HCl solution into a 150-mL beaker. 3. Use a volumetric pipette to measure out 25.00 mL of your unknown acid solution. (Instructor will demonstrate the use of the bulb and pipette). Place the acid solution in an Erlenmeyer flask. Add 2-3 drops of indicator. (Remember, the phenolphthalein changes from colorless to pink at the endpoint.) Place the flask under the burette, record the initial burette reading, and then add base from the burette until you reach the endpoint, as before. Record the final burette reading, discard the sample and repeat until you have at least three successful trials. Return the cleaned burette. EXPERIMENT 17 REPORT Name TITRATION Section Part A: Data Table: (Be sure to record the correct number of significant figures.) A vial or small container with KHP may be provided, or use the chemical directly from the reagent bottle as directed. Trial 1 Trial 2 Trial 3 Initial mass of flask g g g Mass of flask with KHP g g g Initial burette reading of NaOH mL mL mL Final burette reading of NaOH mL mL mL Calculations: Show a complete calculation of moles and molarity from Trial 1 in the spaces marked with *. (Calculations for trials 2 and 3 can be shown in the notebook) Mass of KHP sample titrated g g g Volume of NaOH solution used mL mL mL Convert the mL to Liters for each run L L L Moles of KHP used in each sample* mol mol mol Significant figures are critically important Moles of NaOH reacting mol mol mol Molarity of NaOH solution (calculations)* M M M Average molarity of the standardized NaOH solution M This value will be used in calculations for part B of the experiment. Part B: Data Table: (Be sure to record the correct number of significant figures.) HCl solution Unknown Code Trial 1 Trial 2 Trial 3 Volume of HCl solution used mL mL mL Initial burette reading of NaOH mL mL mL Final burette reading of NaOH mL mL mL Calculations: Show a complete calculation from Trial 1 for each of the following, * then complete the work for the remaining calculations in your notebook. Volume of NaOH solution used mL mL mL Convert the mL values to liters L L L Moles of NaOH used in each sample (Use the molarity of NaOH from part A) mol mol mol * Moles of HCl reacting mol mol mol (Refer to the first equation, in the introduction) Molarity of HCl solution M M M * Average molarity (experimental value) of the unknown HCl solution M (To be filled in by the instructor) Actual molarity of the unknown HCl M Percent error % Name QUESTIONS AND PROBLEMS Section 1. A student weighed out 1.54 g of KHP. How many moles is this? How many moles of NaOH will react with this KHP sample? 2. The equation for the reaction of KHP and NaOH shows a 1:1 ratio for the two reactants. Experimentally, how do you know that the base solution that you delivered from the burette had just as many moles as were in the KHP sample that you weighed out in the beginning of the experiment? 3. What might happen to your calculated NaOH molarity if you use tap water rather than deionized (purified) water to dissolve the KHP crystals or to rinse down the walls of the flask during the titration? (Tap water contains some calcium bicarbonate.) 4. Write and balance the equation for a neutralization of a sulfuric acid solution of unknown concentration by sodium hydroxide. Calculate the molarity of an unknown sulfuric acid solution if a 25.0-mL sample of the acid solution consumes 27.2 mL of 0.138 M NaOH solution in a titration. .

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