Titration of a Diprotic Acid: Identifying an UnknownAj Klatch Kyle Tisi Mr. Edmondson Ap Chemistry 4/19/022/31I.AbstractThe following titration lab emphasizes several concepts important to solution chemistry and acid-base equilibrium. Titrations are careful procedures used in chemistry to determine the concentration of a particular solution. A solution containing a known concentration of base, called the standard solution, is slowly added to the acid until a neutralization reaction occurs between the two. The point at which stoichiometrically equivalent quantities are brought together is called the equivalence point of the titration.This lab involved the titration of solid maleic acid, a diprotic acid, with 0.25M NaOH base. The molecular formula for maleic acid is H2C4H2O4. A diprotic acid differs from a conventional monoprotic acid because it has two ionizable H atoms instead of just one. Each of the H atoms in maleic acid ionize in successive steps. Specifically, maleic acid dissociates in water in two stages:H2C4H2O4 H C4H2O4 + H3OH C4H2O4 C4H2O4 + H3OThe acid dissociation constants for these equilibrium stages are labeled Ka1 and Ka2, respectively. The numbers on the constants refer to the particular proton of the acid that is ionizing. It is important to know that Ka2 is much smaller than Ka1. On the basis of electrostatic attractions, we would expect a positively charged proton to be lost more readily from a neutral H2C4H2O4 molecule instead of a H C4H2O4 molecule. It is always easier to remove the first proton from a diprotic acid than the second proton. Because Ka1 is so much larger than Ka2, almost all the H protons come from the first ionization stage.In this lab, my partner and I performed steps comparable to any acid-base titration. Our first step was to set up the apparatus and make sure that everything was working properly. For example, we rinsed out all of the glassware we were going to be using, including the burette and beakers. We made sure that the burette was properly calibrated and also checked that the stopcock was working correctly. Because we were using a solid diprotic acid (maleic), our next step was to weigh out the designated .120 grams of the solid acid on an electronic balance. It is important to know that we didnt use an acid-base indicator in this titration, but rather a pH probe with an amplifier and a hot plate equipped with a stirring feature. The use of the hot plate and magnetic stirrer eliminated the need to swirl the beaker as the base is added to it. It was important to calibrate the pH probe before beginning any data collection. To calibrate this device, we attached the sensor to a computer port and placed it in 3 different solutions, each of which had a different pH value (4,7, and 10). In the first step of the calibration, the probe was placed in a buffer solution of pH 4.00. From here, we recorded the data on a computer and rinsed off the probe in distilled water. The probe was next placed in a buffer solution of pH 10.00. Again, we recorded the values on the computer and rinsed off the probe, allowing it to stand in a neutral pH 7.00 solution also provided. It was important for us to keep the pH sensor always in a solution to avoid damage.2After all the equipment was set up and running, my lab partner and I now prepared our 100 ml acid solution with the use of a 100 ml volumetric flask. First, we added the weighed out 0.120 grams of maleic acid to a beaker and added about 50 ml of distilled water to completely dissolve the solid. We then poured this solution into the given volumetric flask. We finished preparing this solution by adding remaining distilled water to graduated mark on the 100 ml flask. A 100 ml solution of maleic acid of 0.01 Molar had now been prepared. Next, we added 100 ml of 0.25M NaOH to the zero mark of a 100 ml burette calibrated at +/- 0.2 ml. To a 250 ml beaker, we added the 100 ml maleic acid solution. To this we also placed the magnetic stirrer. Once all of the apparatus was set up, we began the titration.Making sure that the pH computer program was ready, we slowly opened the stopcock and added a steady stream of ONLY drops of NaOH base to the maleic acid beaker solution. When performing this titration, we added drops of NaOH until the pH of the beaker solution increased by a factor of 0.20. After each successive 0.20 pH increase, we recorded the volume of NaOH added to the beaker by highlighting the KEEP icon on the computer screen. All volumes were recorded in mls, and this procedure continued until the pH of the solution in the beaker reached exactly 3.5. Once 3.5 was attained, we changed to 2 drop increments until a pH of 4.5 was reached. We stored the data by again using the KEEP icon on the computer. At pH of 4.5, my lab partner and I went back to the original 0.20 pH increments of change. This continued to a pH of 7.5. After each increment, the burette reading was stored. At 7