Lab S

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Lab S. Acids and Bases: Titrations Jayden Podmoroff 17122912 Partner: Miller Lab performed: Feb 15, 2019 Chem 123 L22 TA: Sydney Fearnley Lab Submitted: Feb 29, 2019 1

Introduction: The purpose of this experiment was to perform different acid-base titrations and observe the use of indicators in the titrations. Conducting this experiment will further the development skills in plotting graphical data and graphical interpretation. A titration is defined as a technique that uses a solution with a known concentration is used to find the unknown concentration of another solution. The titrations involved in Lab S are all acid-base titrations, where two both a strong (HCl) and weak (CH 3 COOH) acid are titrated with a basic solution (NaOH) of known concentration. The chemical reaction of the titration varies depending on whether the acid and base involved is strong or weak. The reaction for strong acid and strong base titrations are given in reaction (I) while the reaction for weak acid and strong base titrations are given in reaction (II). 1 The corresponding reactions are given below. 1 H 3 O + + OH -  2H 2 O (I) HA + OH -  A - + H 2 O (II) Once the acid has been fully neutralized by the base, the unknown concentration can be found using the total volume of base added and the balanced chemical equation reaction of the titration. The data obtained from a titration can be plotted as a graph into what is called a titration curve, where pH is plotted as a function of volume of titrant added. The equivalence point can be found using a titration curve graph by finding the average of the end point and start point, which is asked of Part A (intersecting points between tangent lines and slope of pH jump). 2 Part B seeks to find the endpoint of different titrations with the use of indicators. Indicators are weak acids or bases that are used to monitor the progress of a titration through colour change at certain pH 2

levels. The appropriate indicator must be used for the appropriate titration based off pH range in order to yield accurate results. The change in colour occurs when the indicator changes into its conjugate acid or base form through removal of a proton, using the equation (III) below. 3 HIn + OH -  H 2 O + In - (III) Colour I Colour II Procedure: The procedure of Lab S given by the First Year Chemistry Lab Manual 4 was followed as described below. Part A of Lab S required the use of a pH meter to find the equivalence point. First, the pH meter was calibrated by following steps described in the lab manual. A burette was attached to a retort stand and filled with an appropriate amount of standardized NaOH solution with concentration 0.105 M. 10.00 mL of the hydrochloric acid solution was pipetted into a 100 mL beaker along with about 20 mL of distilled water. The beaker was placed on a magnetic stirrer and a magnetic stirring flea was placed inside the beaker. The initial pH of the solution was recorded using the pH meter. The standardized NaOH solution was titrated very carefully with the HCl solution. The NaOH was added in small increments ranging from 1 mL to 0.1 mL and the reading from the pH meter was recorded after each corresponding addition. Smaller increments were used when finding the equivalence point of the titration and larger increments were used nearing the end of the titration. After completion of the titration, the exact same steps were repeated with the exception of using a acetic acid solution instead of an HCl solution. 5.00 mL of acetic acid was pipetted into a 100 mL beaker along with about 20 mL of distilled water to create the solution. 3

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Part B of the lab required the use of indicators to find the equivalence point of different titrations. First, the provided HCl was titrated with the NaOH solution three times; each time using 3 drops of either bromophenol blue, bromothymol blue, or phenolphthalein. 10.00 mL of the provided HCl was pipetted into a 125 mL Erlenmeyer flask along with about 20 mL of distilled water. The end points were determined with the different indicators and the volumes of NaOH used were recorded. The above steps were repeated again, but this time using and pipetting 5.00 mL of acetic acid as opposed to 10.00 mL of HCl. Three titrations with the three different indicators were performed, each time recording the volume of NaOH used, thus a total of six titrations were performed in part B. Data/Observations: Part A Titration pH at Equivalence Point Volume at Equivalence Point Molarity of Analyte HCl and NaOH 6.97 10.27 mL 0.108 M Acetic Acid and NaOH 8.26 12.67 mL 0.266 M Part B Bromophenol Blue Bromothymol Blue Phenolphthalein Molarity Volume of NaOH added (HCl) 10.4 mL 11.4 mL 10.4 mL 0.120 M Volume of NaOH added (Acetic Acid) 2.3 mL 8.2 mL 13.0 mL 0.273 M Each indicator rapidly changed colour once the solution had reached its pH range. The indicators yielded consistent results in comparison to the pH meter results. It was easy to tell when the inappropriate indicator was used because of the large difference in volume of NaOH used. 4

Discussion: The equivalence point for the HCl titration was found to be 6.97 pH and the equivalence point for the acetic acid titration was found to be 8.26 pH. These results support the fact that the HCl and NaOH titration is a strong acid and strong base titration and that the acetic acid and NaOH titration was a weak acid and strong base titration. This is because an ideal strong acid and strong base titration will always have an equivalence point of close to 7.00 pH and weak acid and strong base titrations will always have an equivalence point >7.00 pH, because when the weak acid is neutralized it is converted into its conjugate base which reacts with water to increase the concentration of H + as seen in reaction (III). 3 Results showed that the ideal indicator for the acetic acid titration was phenolphthalein. This is because Phenolphthalein’s pH range is ideal for the equivalence point of the titration. When analyzing acetic acid’s titration curve, it can be seen that phenolphthalein’s pH range is perfect for its equivalence point while bromthymol blue and bromomethol blue’s ranges are too low on the curve. Unlike the acetic acid titration, all three indicators yielded accurate results for the HCl titration. This is consistent when analyzing HCl’s pH curve because all three indicators’ ranges fall on the equivalence point. Possible sources of error in this experiment would have been related to reading burettes and pipetting liquids. In Part A, the burette containing NaOH was recorded about 25 different times for each titration. The uncertainty in a burette is  0.02 mL, therefore adding that uncertainty for each reading would become a major source of error in the experiment. There is also uncertainty of  0.02 mL when using 5 mL and 10 mL pipettes, which were involved in each titrations process throughout the experiment. Another source of error could have came from the titrations in Part B. When using a burette to titrate, it is very difficult to add the exact volume of liquid needed to reach the end 5

point. There is often a small amount of excess liquid added due to human error, contributing to more uncertainty in the experiment. Conclusion: The concentration of The provided HCl solution was found to be 0.108 M and the concentration of acetic acid was found to be 0.273 M, using the pH meter technique. The most effective indicator for the acetic acid titration was phenolphthalein and the HCl titration yielded accurate results with all three indicators. References: [1] First Year Chemistry Lab Manual: Chem 111/113 & Chem 121/123 , University of British Columbia: Kelowna, BC, 2018-19; p 125 [2] Waser, J. Acid-Base Titration and Distribution Curves , Chem. Educ. , 1967 , 44 (5), p. 274 [3] First Year Chemistry Lab Manual: Chem 111/113 & Chem 121/123 , University of British Columbia: Kelowna, BC, 2018-19; p 126 [4] First Year Chemistry Lab Manual: Chem 111/113 & Chem 121/123 , University of British Columbia: Kelowna, BC, 2018-19; p 127-128 6

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