Formal Lab Report
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Formal Lab Report

Course Number: CHEM 1AL, Fall 2012

College/University: Berkeley

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Stephen Frianeza Chemistry 1AL Lab 100 9/24/12 Acids in the Environment II: Indicator Titration Abstract The task in this experiment was to accurately determine the molarity of an HCl solution of unknown concentration. The solution was diluted and then a titration was performed using the base Tris(hydroxymethyl)aminomethane (TRIS) and bromocresol green indicator. TRIS in solution with the indicator turns blue, and...

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Frianeza Chemistry Stephen 1AL Lab 100 9/24/12 Acids in the Environment II: Indicator Titration Abstract The task in this experiment was to accurately determine the molarity of an HCl solution of unknown concentration. The solution was diluted and then a titration was performed using the base Tris(hydroxymethyl)aminomethane (TRIS) and bromocresol green indicator. TRIS in solution with the indicator turns blue, and the dilute HCl was added until it turned green. After three trials with less than 4% difference were performed, the average molarity of the dilute HCl was calculated to be 0.1123 M. Introduction The chemical reaction taking place in this experiment is the following: HCl(aq) + TRIS + TRISWhere TRIS = (HOCH2)3CNH2 The goal of the experiment is to accurately determine the molarity of the unknown HCl solution to 4 significant figures. The measuring instrument with the least number of significant figures is the buret with 4, so the calculations will be accurate to at most 4 figures. By figuring out the molarity of the solution, the potential danger of it will be known and it can be used to obtain accurate data in other experiments. HCl will be titrated to the basic TRIS solution, and the color shown by the indicator will be carefully monitored for changes. The method of titration was chosen due to the accuracy of measurements; burets are able to measure to 0.01 mL. The theory behind this experiment is that, as an acidic solution is added to a basic solution, the hydronium and hydroxide ions react to form water. This lowers the pH of the basic solution in the flask. At a pH of roughly 5, the equivalence point, bromocresol green turns green, and the number of moles of HCl and TRIS are equal. At this point, the volume of HCl used was measured and, knowing the amount of TRIS used, the number of moles of reactants was calculated, leading to the determination of the molarity of the unknown HCl solution. Methods 5 mL of the concentrated HCl was pipetted into a 250.00 mL volumetric flask, and water was added until the line. The buret was rinsed twice with the dilute HCl. Between 0.15 and 0.25g of the TRIS was measured out and then dissolved into a 250 mL Erlenmeyer flask so that the final solution volume was 50 mL. 5 drops of bromocresol green were added. Titration of the HCl began at a fast rate, swirling as the acid was added, and slowed down as the equivalence point was approached. Upon reaching the equivalence point, when the solution turned green, titration was stopped and final volume of HCl was recorded. The following measurements were taken: mass of TRIS to the nearest 0.0001 grams, and initial and final volume of HCl to the nearest mL. 0.01 The procedure is from the General Chemistry 1A Laboratory Manual, Michelle Douskey, Fall 2012. Results Trial TRIS (g) 1 0.2420 2 0.1776 3 0.1533 Initial HCL (mL) 0.60 0.88 0.82 Final HCl (mL) 18.26 14.01 12.10 Used HCl (mL) 17.66 13.13 11.28 Molarity (M) 0.1131 0.1117 0.1121 Knowing that the molar mass of TRIS is 121.14 g/mol, the following calculations were performed to determine molarity of HCl: For trial 1: The calculations were repeated for Trials 2 and 3. Average Molarity: 0.1123 M All digits from all measurements were recorded and final number of significant figures was determined after all calculations. In order to determine the concentrated molarity, the following calculation was performed: Class average for the molarity: 0.0992 M Percent difference from class average: 13.21% Discussion and Conclusions With unknown concentration of an acid or base, titration is used because the method employs very accurate measurements and it can be performed slow enough that the process can be stopped at any point. One very important observation from this experiment was that just one small drop was able to change the color of the solution from blue to yellow instantly. In all trials performed, the proper green color was not obtained, with each trial turning yellow. Inexperience in executing accurate titrations and being unable to successfully perform partial drops played a large factor in not obtaining the green equivalence point solution color, and the additional HCl added which changed the solution yellow was likely the reason for the high molarity calculated and percent difference from the class average being so great. However, the molarities calculated in all three trials were within 1% of the average of 0.1123 M, so the titration was consistent each time. Overshooting the equivalence point led to believing the molarity of the HCl was higher than it should have been, in this case 0.0131 M higher. Therefore there were fewer moles per liter than the calculations say, so a greater volume of HCl would have been required to reach equivalence. Using the data from trial 1 and the class average for molarity, the amount of volume that should have been used was 20.01 mL, compared to the experimental value of 17.66 mL. Solving for V yields: However, this calculation does not agree with the intuition that in overshooting the equivalence point, a lower volume must have been required to reach equivalence. Thus, there must be other factors involved that led to the calculated average molarity being greater than the class average. Further experience in operating and reading the laboratory equipment should lead to more accurate experiments.

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