Aspirin, also known as acetylsalicylic acid, is a commonly used pain reliever and anti-inflammatory medication. In a titration lab, students can learn about the properties of aspirin and how to determine its purity and concentration.
The titration process involves the precise addition of a known concentration of a solution (called the titrant) to a solution of unknown concentration (called the analyte). The goal is to determine the concentration of the analyte by measuring the volume of titrant required to react with it.
In the aspirin titration lab, students will start by preparing a standard solution of sodium hydroxide (NaOH), which will serve as the titrant. The concentration of this solution should be known with high precision, typically through the use of a calibrated burette or pipette.
Next, students will prepare a solution of aspirin by dissolving a known mass of the medication in a solvent such as water. The concentration of the aspirin solution can then be determined by titrating it with the NaOH solution.
To perform the titration, students will use a piece of equipment called a burette, which allows them to dispense precise volumes of the titrant solution. The volume of NaOH added to the aspirin solution will be recorded at regular intervals, typically using a calibrated pipette or graduated cylinder.
As the titration progresses, the solution will change color as the acidity of the aspirin is neutralized by the alkalinity of the NaOH. The endpoint of the titration, also known as the equivalence point, is reached when the solution has been neutralized and no further change in color is observed.
The volume of NaOH required to reach the equivalence point can be used to calculate the concentration of the aspirin solution. This calculation is typically performed using a formula known as the "titration equation," which takes into account the known concentration of the NaOH solution, the volume of NaOH used, and the molecular weights of the reactants.
In addition to determining the concentration of the aspirin solution, students can also use the titration lab to verify the purity of the medication. If the aspirin is not pure, the volume of NaOH required to reach the equivalence point will be different from what is expected based on the concentration of the aspirin solution. This difference, known as the "titration error," can be used to calculate the percent purity of the aspirin.
Overall, the aspirin titration lab is a valuable educational tool that allows students to learn about the properties of aspirin and how to determine its concentration and purity. Through hands-on experimentation, students can gain a deeper understanding of the titration process and its importance in the field of chemistry.
Lab Report; Titration of Aspirin complianceportal.american.edu
DATA TABLE Part I Synthesis of Aspirin Mass of salicylic acid used g Volume of acetic anhydride used mL Mass of acetic anhydride 1. The number of moles of base consumed and the number of moles of acid in the sample can be calculated from the volume of the base needed to obtain the color change. In order to approximate the molar mass of aspirin the mass of aspirin would be divided by the moles of aspirin present. SAMPLE DATA Melting temperature determination for aspirin sample Organic Chemistry with Vernier 11 Spectrum for salicylic acid with iron III nitrate Absorbance vs. The absorbance of 0. Obtain and prepare a buret for titration by washing and conditioning it with titrant solution NaOH.
Calculate his percent purity. Vapors are strongly irritating. Now that you have a rough idea of the melting temperature, a more accurate determination of the melting temperature can be made. After the sample is completely dissolved, add 90 mL of distilled water, and stir the solution again. After students complete the Part 1 synthesis, they will need to dry the sample. Van 1 Kayla Van Ms.
Quantitatively transfer the solution from the beaker to a 250 mL volumetric flask. Repeat as necessary to assure an accurate and reproducible result. In case of contact with skin, rinse the area with large amounts of water and notify your instructor. To determine the concentration of the salicylic acid impurity in the treated aspirin sample, interpolate along the regression line to convert the absorbance value of the unknown to concentration. Aspirin titration lab Itamar Aviel Department of chemistry and biochemistry, Queens College - CUNY CHEM 1134-11, Fall 2020 Instructor: Babak Gandjian Abstract This lab allowed us to undertsnad how to identify the amount of aspirin found in a tablet through an acid- base titration. The EMM is the mass of sample in mg which produces one mmole of H + ions upon titration.
Determination of Aspirin Concentration by Back Titration
NaOH was slowly added while constantly swirling until the solution turned pink, and its volume was recorded. Disconnect the Melt Station and connect the Spectrophotometer to the USB port of LabQuest or a computer. This might be due to errors which can be determinate or indeterminate, it also depends on how large the error is made. The titre values were noted when the end point was determined that is when the solution turned from pink to clear. For a pure compound with one acid group per molecule, this should be the actual molar mass of the compound.
The following equation, H-Asp aq + NaOH aq Na-Asp aq + H2O l illustrates the neutraliztion reaction that occurred in the experiment. Addition of denaturant makes the product poisonous. The rest of the calculations are normally done. Background Information Titration helps determine the concentration of a solution. Transfer the contents of the flask to a Büchner funnel assembly. Moderate fire risk flash point 49°C.
Acetic anhydride: Strongly irritating and corrosive. You are now ready to collect data for the five standard solutions. The point in the titration where enough of the titrant has been added to react exactly with the analyte is called the equivalence point, and occurs when moles of titrant equals moles of analyte according to the balanced equation. If an indicator is added into the reaction, a color change will take place when an acid will completely neutralize, and on the drop of the excess base would be added. Your two primary objectives in this experiment will be to synthesize and analyze aspirin.
The search for a milder form of this pain reliever led to the successful synthesis of acetylsalicylic acid by the German chemist Felix Hoffmann in 1893. Measure and record the absorbance value of the treated aspirin sample. The blue light will turn on indicating that the fan is cooling the Melt Station. Causes severe eye damage. Wear goggles at all times in the chemistry laboratory. What is the theoretical yield of aspirin in your synthesis? On the other hand, the primary standard is the KHP.
Steps 22—25 will guide you through the set up and testing of your aspirin sample. Waste disposal containers were located by the hood, so the chemicals were disposed of carefully based on the. Carefully observe the temperature vs. Primary standards are acids that completely dissolve in water, have high molar mass, and are not hygroscopic. Wipe the outside, place it in the device, and close the lid.
After about 15 seconds, turn the suction back on. Record the precise value in your notebook and on the data sheet. When most of the liquid has been drawn through the funnel, turn off the suction and wash the crystals with 5 mL of cold, distilled water. Determinate errors might be constant or proportional having measurable and fixed value which can be corrected and as the experiment proceeds, error increases with the magnitude of the measurement. Part II: Determination of the purity of aspirin tablets Second, you will titrate a sample of your aspirin acetylsalicylic acid with the standardized NaOH to determine the moles of acid in a given weight of your product. Add two or three drops of phenolphthalein indicator to each solution and titrate with the standardized sodium hydroxide solution provided. Because the mole ratio is 1:1, the number of moles of the base will equal the number of moles of acid at the equivalence point of the titration.
Assuming the aspirin is not contaminated with other acids, the titration allows you to quantitatively determine the purity of your aspirin. A titration process can be used to determine ASA content in commercial aspirin tablets. Transfer 5 mL of the aspirin solution from the 250 mL volumetric flask to a clean and dry 100 mL volumetric flask. Determine the optimal wavelength for creating the standard curve and set up the mode of data collection. Error is particularly high.