Acid Base Laboratory
2.04 g of KHP, 100ml volumetric flask, distilled H2O, approximately 0.1 M of NaOH, Vinegar, Phenolphthalein, 250ml Erlenmeyer flask, weighing balance, Graduating Cylinder, burette and pH meter were used in our experiment.
In our first part of our experiment to prepare a primary standard, 0.1 M solution of KHP, we carefully weighed out 2.04g of KHP in a weigh paper using the weighing balance and transferred the measured KHP into a 100ml volumetric flask. We rinsed the weigh boat with distilled H2O to make sure that all the measured KHP is transferred into the flask and added distilled H2O into the flask until the meniscus of the water touches the line of the flask. We did not measure the amount of H2O added since water doesn’t involve in the reaction. We then swirled and inverted the flask until all the KHP is dissolved into the H2O. KHP+H2O-------H3O+ + KP-
In our second part of our experiment to determine the molarity of NaOH, we rinsed a burette with about 5ml of water and another 5ml of NaOH to prepare it for the experiment. We then measured exactly 25ml of the prepared KHP from the previous part into a 250ml Erlenmeyer flask and added 2 drops of the indicator, phenolphthalein. Next we filled the burette with the approximately 0.1 M NaOH so that our initial volume will be 0.0mls. We then started our titration by carefully pouring a little amount of the base into the prepared 25ml of KHP with phenolphthalein until we reached the end point. Since the molarity of the NaOH was designed to be more or less the same with the acid, we knew that it would take equal amount of the base to reach the endpoint. Thus we added the first 20 mls in about 5ml increment and swirled to mix after each addition then added just 6-7 drops for the next mls. After we started observing a pink color, we started adding just 2 drops at a time and swirl the flask. When we reached a persistent pink color, we recorded that final burette to the tenth place as the end point (equivalent point). After that, we calculated the molarity of the NaOH based on the reaction between the acid and the base and the prepared molarity and volume of the acid and the volume reading of the base. We then repeated the same experiment to make sure that our calculated molarity of the base are within 10% difference.
In our third part of the experiment to determine the molarity of vinegar, we measured exactly 2.5ml of vinegar using a graduating cylinder into a clean non-dry Erlenmeyer flask. We then added 2 drops of phenolphthalein. Next, we refilled the burette with the base, NaOH and recorded the initial volume as 0.0ml. We then titrated with NaOH very carefully in a 1ml interval since we don’t know the expected equivalent point. Once the color begins to persist, we slowed our titration to 2-3 drops at a time so that we wouldn’t overshoot the endpoint. We then recorded the final volume when we observed a permanent pink color at the equivalent point and calculated the molarity of the acid using the formula M acetic acid=Moles NaOH/Volume in liters of the acid (2.5 ml) We then had to repeat the experiment twice to make sure our molarities agree within 10% of the values.
In our fourth part of the experiment to determine the equilibrium constant of vinegar, we measured 2.5mls of vinegar into a wet flask. Unlike the previous part of the experiment, we used a pH meter instead of a phenolphthalein indicator. We added 15ml of distilled water so that we could have enough liquid to dip the electrode/pH meter and recorded the initial pH of the vinegar before the titration started. Then, we filled out the burette and recorded the initial volume as 0. Next, we added 0.5ml of the NaOH into the vinegar flask and used a pH meter to record the volume. We added the base in a small increment to observe the shape of the titration curve. We repeated the step until we reached...