Titration of Aspirin Tablets In this lab, you will determine ✓ Solved

In this lab, you will determine the amount of acetylsalicylic acid (2-ethanoyloxybenzene carboxylic acid) in a batch of aspirin tablets through a titration simulation. You will do a simulation that guides you through the process step by step learning a variety of techniques. This is a common practice in industries to test the purity of the samples and to see if the amount of acetylsalicylic acid matches the amount on the box.

This is called Quality Assurance, random samples from each batch of tablets are tested in a variety of ways to see if they fall within the conditions for use and sale. To find the amount of acetylsalicylic acid in the tablets you will be performing a virtual acid-base titration.

An acid and a base react to produce a salt and water by transferring a proton (H+). The reaction occurring in your titration involves Aspirin (C9H8O4) and Hydroxide. Aspirin is a weak acid, which indicates that it dissociates and ionizes in solution.

Before you begin the simulation, we will go through some basics about titrations and technique. A titration is a procedure for determining the concentration of a solution (analyte) by allowing a carefully measured volume of this solution to react with another solution whose concentration is known (the titrant). The equivalence point occurs when moles of titrant equals moles of analyte according to the balanced equation.

In this lab you will be performing an acid-base titration. Follow the instructions to complete the lab simulation, including preparation of the aspirin tablets and conducting the titration through the virtual simulation.

After performing the three titrations, calculate the moles of NaOH used and the grams of aspirin in the tablet for each trial, and analyze the quality and purity of the aspirin tablets. Address potential sources of error and reflect on your results in the post-lab questions provided.

Paper For Above Instructions

The determination of the amount of acetylsalicylic acid in aspirin tablets through titration is an essential practice not only for ensuring quality control in pharmaceutical production but also for understanding fundamental acid-base chemistry. Aspirin, or acetylsalicylic acid, serves as a common medication used worldwide, and accurately knowing the active ingredient's quantity is crucial for dosages and safety.

In performing the titration simulation, we follow a structured method that underpins analytical chemistry. The titration process begins by preparing the aspirin tablets where ten tablets are crushed using a mortar and pestle. This step is crucial because it ensures that the acetylsalicylic acid is fully exposed and can adequately react with the titrant.

Next, we prepare the analyte, combining the crushed aspirin with a solvent to facilitate the reaction with sodium hydroxide (NaOH), a strong base acting as the titrant. Rinsing the glassware, especially the graduated pipette and the buret, is of utmost importance. This is to prevent contamination that could skew results and impact the reliability of the measured volumes.

During the titration, the carefully metered addition of NaOH to the analyte solution leads us to the equivalence point, which signifies the total neutralization of the acetylsalicylic acid by the hydroxide ions from NaOH. Monitoring the pH with a probe allows us to observe changes in the acidity of the solution, identifying when the reaction nears completion. The use of phenolphthalein as an indicator is especially effective since it shows distinct color changes, turning pink as the endpoint is approached.

It is crucial to document all readings meticulously, including the initial and final volumes of NaOH dispensed throughout the three titrations. Such data will aid in calculating the moles of NaOH reacted as well as the corresponding moles and mass of acetylsalicylic acid using stoichiometric principles. Stoichiometry plays a significant role here as it converts the volume of NaOH used into moles through the molarity (concentration expressed in moles per liter), which is then used to calculate grams of the aspirin present in the tablet.

The theoretical molar mass of acetylsalicylic acid (C9H8O4) is approximately 180.16 g/mol, and using this value, we can compute how many grams were present based on the moles derived from our titration results. After performing calculations for each trial, we can determine the average mass of aspirin per tablet and compare this to the expected range provided by the manufacturer, which states that Bayer Aspirin contains approximately 325 mg of aspirin.

Analyzing the results may reveal discrepancies such as obtained averages falling outside the expected range of 310 mg to 315 mg of acetylsalicylic acid. Possible sources of error could include (but are not limited to) improper titrant measurement, inaccuracies in the dissolution of the tablets, or human errors in visualizing the endpoint of the titration.

In the post-lab questions, explaining why phenolphthalein is a superior choice over other indicators like bromothymol blue, particularly for a weak acid-strong base titration, would typically center around its more distinct and visible color change that aligns closely with the pH levels around the equivalence point. Understanding these concepts enhances our ability to perform manual titrations in practical lab settings in the future.

Furthermore, providing a reflection on the necessity of rinsing glassware before use emphasizes the importance of preventing cross-contamination, which is vital in obtaining accurate and reproducible results. This underlines the critical precautions that must always be observed in laboratory settings.

In conclusion, the titration of aspirin tablets not only serves as a practical investigation into acid-base chemistry, but it also exemplifies methods of quality assurance in pharmaceuticals, ensuring that consumers receive medication that is effective and safe. Accurately determining the mass of the active ingredient is vital for both public health and scientific integrity.

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