TITRATION OF AN ACID WITH A BASE 1 NOTE: You are required to view the podcast entitled Use of Burets for Titrations before coming to lab this week. To view the podcast, consisting of eight episodes, go to http://podcast.montgomerycollege.edu/podcast.php?forcemp4play=yes&rcdid=165. You may have prelab quiz questions that are directly from the podcast. If your instructor believes that you have not watched the podcast, you may be sent home with a grade of zero! INTRODUCTION In this experiment we will learn how to prepare aqueous solutions and calculate their concentrations (in terms of molarity). An acid solution whose concentration is known with great precision will be made. Then, this solution will be reacted very carefully with a base solution whose concentration is known only approximately. By calculation, we will then determine the base solution's exact molarity. The acid-base reaction will be carried out as carefully as possible. The goal is to combine the exact stoichiometric amounts. In other words, we will try not to have either reactant be in excess. Such a careful addition of one reactant with another is called a titration. We will titrate using instruments called burets. In reality, there may be a slight excess of one reactant, but we try to keep this excess as low as possible. The point at which the stoichiometrically precise amounts have been combined is called the equivalence point. Once we have determined the exact molarity of our base solution, it will be saved and used in the next lab. Any solution whose concentration is known with great precision (usually four or more significant figures) is often referred to as a standardized solution. In this lab, we will first make an "standard" acid solution, which will then be used to standardize the base solution. The reason we do not make a standard base solution to start with is that the most common bases, such as NaOH and KOH, cannot be weighed out on a balance without absorbing some moisture from the air. Therefore, the mass of base cannot be known with great certainty. On the other hand, there are many solid acids which do not absorb moisture from the air in the process of weighing on a balance. The acid we will use is oxalic acid. The solid form of it is a hydrate, specifically H2C2O4 2H2O. The waters of hydration must be included when calculating the molar mass of the solid; however, once dissolved in liquid H2O, the waters of hydration are no longer considered. The base we will use is NaOH. The balanced equation for the reaction that will occur during the titration is: H2C2O4(aq) + 2 NaOH(aq) Na2C2O4(aq) + 2 H2O(l) According to this reaction, one mole of oxalic acid reacts with two moles of sodium hydroxide to produce one mole of sodium oxalate and two moles of water. Since there is no visual evidence that this reaction is happening, we will add an acid-base indicator and watch for its color to change to signal the end of the titration. In theory, this "end point" should be the equivalence point.
2 PROCEDURE A. Preparation Of A Base Solution Whose Concentration Is Approximately Known 1. Calculate the volume of 6 M NaOH needed to prepare 500 ml of a solution with a concentration of 0.2 M. (See prelab question 1). Confirm the amount you have calculated with your instructor. 2. With your graduated cylinder, measure this much 6 M NaOH and pour it into a 500-mL volumetric flask. ALWAYS HANDLE NaOH SOLUTIONS WITH CARE. Since you will eventually be standardizing this solution by titration, it is not a problem if you add slightly more or less of either solute or solvent. 3. Fill the flask with deionized water until the meniscus sits on the line inscribed on the neck of the flask. Stopper the flask and shake thoroughly (to assure an even distribution of solvent). 4. Label the flask and set this solution aside. B. Preparation Of An Acid Solution Whose Concentration Is Well Known 1. Calculate the mass of oxalic acid dihydrate (H2C2O4 2H2O) needed to prepare 250 ml of a 0.1 M H2C2O4 solution. (See prelab question 2b.). Confirm the amount you have calculated with your instructor. 2. NOTE - SOLID OXALIC ACID IS A SKIN IRRITANT. IF YOU ACCIDENTALLY GET SOME ON YOUR HANDS, WASH IMMEDIATELY. DO NOT RUB YOUR EYES, NOSE OR MOUTH. Weigh out this amount of solid carefully on an analytical balance. It is not necessary to weigh out precisely the same mass of oxalic acid that you calculated above. However, you must know exactly how much you did weigh out. This mass will serve as the basis for the calculation of the molarity of the oxalic acid solution. 3. Carefully, without spilling, transfer the oxalic acid to a clean 250-mL volumetric flask. 4. Add deionized water until the flask is about half full and swirl until all the solid dissolves. Then fill the volumetric flask with more deionized water, using an eyedropper pipette for the last milliliter or so, until the meniscus sits on the line marked on the neck of the flask. 5. Stopper the flask, turn it upside down, shake it vigorously and turn it right side up. Repeat this shaking at least ten times to be certain that the mixture is homogenous (and therefore, a true solution). 6. The volume of this solution is 250.0 ml. Calculate its molarity before going on. C. Standardization of the NaOH Solution 1. Clean two burets by the following procedure. a) Add about 5 ml of your base solution and add it to one of the rinsed burets, swirling so that all of the buret surfaces, including the tip, come in contact with the solution; pour out this base. Repeat this twice more. By the end of the NaOH rinsing, the solution wetting the buret will be the same as the base solution being standardized.
b) Repeat the procedure in part b with three small portions (no more than 5 ml each) of standard acid for the other buret. 2. Fill each buret to below the top line on the buret with the respective solutions, making certain that the tip is full of solution and free from bubbles. Label both burets. Now record the starting volume of each buret. Be sure to read the buret from top to bottom. Record buret readings to the nearest 0.01 ml. 3. Add approximately 20-25 ml of acid to a 125-mL Erlenmeyer flask from the acid buret. (You will record the final volume later.) Touch the tip of the buret to a dry side of the flask so that any liquid on the buret tip will be transferred to the flask. Rinse the sides of the Erlenmeyer flask with a small amount of deionized water from a squirt bottle. This process is called "tipping off." Tipping off is essential every time liquid is dispensed from a buret to another container. While tipping off, do not get any water on the buret tip. (Some students question the step of adding deionized water. It does change the concentration of the solution being used. However, it does not change the quantity of solute in the flask. Since the reaction we will conduct is between the acid and the base, any additional water in the reaction flask will not have any effect on the actual reaction.) 4. Add 5 drops of phenolphthalein to the acid solution in the Erlenmeyer flask. 5. With constant swirling of the flask, carefully add the NaOH solution from the base buret to the Erlenmeyer flask until the solution in the flask remains a uniform pink for at least 30 seconds. When a pale pink color persists, you have reached the end point of the titration. You may add the base rapidly at first, then more slowly and, near the end point, drop by individual drop. If you overshoot the end point by adding too much base, the Erlenmeyer contents will be a dark pink color. This is not a disaster. Add more acid drop-wise until the color disappears and then re-add the base drop-wise until a pale pink endpoint is achieved. This recovery from a too-dark endpoint is called "back-titrating." Back-titrating is possible in this lab, but it is not always possible, so you should learn to approach your endpoints carefully. 6. Record the final volumes for both the acid and base burets. The difference between the initial and final readings is the volume of the solution added to the Erlenmeyer flask. 7. Perform at least one more titration (steps 3-6 in section C), making sure that the burets do not become empty in the middle of a titration. Refill the buret if it appears that it may not contain enough solution for the entire titration. 8. You should perform as many titrations as it takes until you get two with consistent results. Check your data with your instructor. 3 D. Finishing and Cleaning 1. When you are finished, save and label your aqueous NaOH solution as directed by your instructor. The label should have your name, contents of the container and instructor's name. 2. Clean the burets and store them upside down, with the stopcocks open.
4 TITRATION OF AN ACID WITH A BASE Data and Calculations Name Partner A. Preparation Of The Base Solution volume of 6 M NaOH used ml volume of solution 500.0 ml molarity of base solution (one sig. fig.) M (this is an approximate molarity, and is NOT to be used in the tables in part C. on page 5) B. Preparation Of The Standard Acid Solution mass of H2C2O4 2H2O used g moles of H2C2O4 2H2O used mol volume of solution 250.0 ml molarity of oxalic acid solution (4 sig. figs.) M (this is a precisely-known molarity) Show your calculations for all of the above quantities in the space below.
5 C. Standardization of Sodium Hydroxide Solution SHOW ALL CALCULATIONS BELOW ON A SEPARATE SHEET OF PAPER! Acid Buret Titration 1 2 3 4 5 final buret reading ml ml ml ml ml initial buret reading ml ml ml ml ml volume of acid used ml ml ml ml ml Base Buret Titration 1 2 3 4 5 final buret reading ml ml ml ml ml initial buret reading ml ml ml ml ml volume of base used ml ml ml ml ml Molarity Titration 1 2 3 4 5 moles of acid used mol mol mol mol mol moles of base used mol mol mol mol mol volume of base soln* used (in L) L L L L L molarity of base M M M M M * soln = solution average molarity of base solution
6 TITRATION OF AN ACID WITH A BASE Post-lab assignment Name Partner 1. A student performs a titration that requires 40.00 ml of 0.5000 M KOH to completely react with 30.00 ml of a H3PO4 solution. What is the molarity of the H3PO4 solution, if the two compounds react according to this equation? (HINT: DO NOT use M1V1=M2V2 to solve for the molarity of acid involved in a titration!) H3PO4 (aq)+ 3 KOH(aq) K3PO4 (aq) + 3 H2O(l)
7 TITRATION OF AN ACID WITH A BASE Pre-lab assignment Name 1. In part A of the procedure, you will make a 0.2 M NaOH solution by diluting 6 M NaOH. How many milliliters of 6 M NaOH are required to make 500.0 ml of 0.2 M NaOH? Even though it will be breaking significant figure rules, give the answer to the nearest tenth. 2. In part B of the procedure you will be dissolving oxalic acid dihydrate in water to make 250.0 ml of an acid solution with a molarity of 0.1 M. a) Calculate the molar mass of H2C2O4 2H2O to five significant figures. b) Calculate the mass of oxalic acid dihydrate required in order to make 250 ml of a solution whose molarity is 0.1 M. Again, go ahead and break significant figure rules to give the answer with 5 significant figures.
8 3. A student performs a titration according to the procedure of this lab. During the process, 15.69 ml of a 0.3059 M H2C2O4 solution are added to a 24.10 ml sample of a NaOH solution with an unknown molarity. The goal is to calculate the NaOH solution s molarity. Do so by the following steps. This is a similar process required for the calculations section of the lab. a) Determine the moles of H2C2O4 used in this reaction (use the molarity equation). b) Determine the moles of NaOH used in this reaction (consult the balanced equation for the reaction). c) Convert the volume of NaOH solution from milliliters to liters. d) Now that you know the number of moles of NaOH in this certain volume of solution, you have enough information to calculate what the molarity of the NaOH solution was before the reaction (and therefore, the molarity of the remaining, unused NaOH solution). Last revised 11/14/2016 DN