Three drops of phenolphthalein indicator was added into the HCl solution. The stopcock was opened and the NaOH solution was added into the HCl solution. The flask was swirled to mix the solutions and titrate to a faint pink end point. Recorded the final volume on the butret and used the final volume as a beginning volume to repeat the titration. At the second titration, the experiment was exactly the same as the first titration but the H2SO4 solution was used to instead of the HCl solution.
Repeat steps 1-6 until three trials have been recorded. Results: Trial | Initial Volume | Final Volume | Total Volume | Color of the Solution | pH | 1 | .60mL | 8.6mL | 8mL | Bright Pink | N/A | 2 | 8.6mL | 14.9mL | 6.3mL | Semi-BrightPink | N/A | 3 | 14.9mL | 21.3mL | 6.3mL | Pale Pink | 7 Yellow | Analysis: The data that we collected, made us infer that the indicator, phenolphthalein, changes to a bright pink color in the presence of a basic solution. Our data also showed that the neutral solution had a lesser amount of NaOH than HCl. This would mean that the NaOH was more concentrated than the HCl. We then calculated the molarity of NaOH with the volume found in our data.
=n (mol)/v (dm3) Whereby c=concentration, n=number of moles and v=volume used. (JOHN GREEN AND SADRU DAMJI, PG 27.THIRD EDITION) Volume=0.1dm3 Concentration of acid =0.0048996mol/0.1dm3 =0.049moldm⁻3 Also, to calculate for the average titre volume of the acid used, I summed up the three consistent values from my experiment and divided it by 3(sarps series, 2009) thus Average Titre volume= (19.60+19.60+19.60) 3 =19.60cm3 Also, to calculate for the concentration of sodium hydroxide, I used The dilution factor, C₁v₁=C₂v₂ (JOHN GREEN AND SADRU DAMJI pg28, third edition) whereby C₁=concentration of the acid used=0.049moldm⁻3 v₁=volume of acid used (titre volume) =19.60cm3 C₂=concentration of the base, v₂=volume of the base used=10cm3 Therefore C₂=
Date: 6/3/2012 Title : Synthesis of an Active Pharmaceutical Ingredient, Aspirin ( Including Raw Materials and Product Purity Checks ) Name: Class: Lab partners: Aim: The aim of this experiment is to prepare and recrystallise Aspirin, determine the percentage yield by mass of product, preform a test to check for phenols, and characterise it by its melting point. The following is the reactions scheme: Method: -[Lab Manual Pages 3, 4 & 5]- Results and Calculations: Compound Salicylic Acid Acetic Anhydride Aspirin Appearance White Powder Clear, Colourless liquid White Crystals Molar Mass (g/mol) 138g/mol 102g/mol 180g/mol Density if liquid (g/cm3) 1.08g/cm3 Weight used (g) 2.868g 5.4g N/A Moles used (mol) 0.0207mol 0.0529mol N/A Limiting Reagent N/A Theoretical Yield of Products in moles N/A N/A 0.0207M Theoretical Yield of Product in Grams N/A N/A 3.726g • Calculations from table ^ : Moles of Salicylic Acid : Moles= Mass X Mr => M = 2.868 X 138 = 0.0207moles Moles of Acetic Anhydride : M = 5.4 X 102 = 0.0529moles Theoretical Yield in moles : moles of limiting reagent = 0.0207moles Theoretical Yield in Grams : Moles of limiting reagent X Mr of Product : = 0.0207mol X 180g/mol = 3.726g • Melting Points Compound Salicylic acid (raw material) Recrystallised Aspirin Expected MP Value 157-161oC 138-140oC Actual MP Value 158-160oC 137-139oC • Percentage yield: Percentage yield = 2.214/3.72 ₓ 100 = 59.52% => 60% • Test for phenols: Samples Crude Aspirin Recrystallised Aspirin Salicylic Acid Colour change of Samples + 1cmoC methanol Yellowish Yellowish Purple Inference Phenol not present Phenol not present Phenol Present Post- Practical Questions : Q1 What did your melting point indicate about the purity of your product? Explain your answer.
Identify that your Excedrin is a mixture of organic molecules using thin layer chromatography (TLC), melting point (mp) and Proton Nuclear Magnetic Resonance Spectroscopy (1H-NMR). 2. Determine the solubility of the components of Excedrin in various solvents. 3. Separate the components of your poisoned Excedrin using solubility characteristics and extractions.
The extraction process is when a solvent, dichloromethane (15mL) is added to the filtrate in a separatory funnel; the mixture is gently swirled together 3 times, and stopcock is released in between to vent the funnel. Dichloromethane (including the emulsion) is then drained from the bottom into a 50mL Erlenmeyer flask. Same extraction process is repeated on the same filtrate and the dichloromethane is, once again, let out to the same 50mL Erlenmeyer flask as before. The combined dichloromethane solution and water (20mL) is poured into a rinsed separatory funnel. Mixture is gently swirled and drained out into an Erlenmeyer flask.
The third test will utilize thin layer chromatography to evaluate the purity of the aspirin as well as testing for the presence of leftover salicylic acid or other by products of the reactions. Experimental: Week 1: For the synthesis of the aspirin, 250 mL of water was boiled. 1.5 g. of salicylic acid were poured on a test tube. Then, 3.5 mL of acetic anhydride and four drops of 85% phosphoric acid were added. A cotton ball was placed to prevent vapor escape.
The Solubility of Calcium Hydroxide Introduction: Experiment 16 was done to get an understanding of the solubility of calcium hydroxide, or the slightly soluble salts, gather the knowledge to calculate the Ksp using a primary standard for determining the concentration of an acid, and learning the knowledge to prepare a saturated solution of a slightly soluble salt. This experiment has a few different parts to complete to gather all the information needed to get the experimental values that are seen in the text books used every day. Those values were at some point one by one experimentally calculated as in this experiment. For the first part of this experiment is –saturated solution, for each pair of students should prepare the saturated Ca(OH)2 solution. Of course the saturated solutions take time this experiment is done in a two day lab setting part one should be completed during the first day and ready for use on day two.
Antacid Tablet Analysis: An Experimental Approach to Acid-Base Titration Erin Votaw, Alioune Sakho, Matthew Wahhab, Alex Williamson, Richard Vela (Group 6) General Chemistry 102 (CHEM 102), ISP SCUHS February 12, 2015 Abstract: In this laboratory experiment, we sought to determine the active component of an antacid tablet. We hypothesized that if we correctly executed the acid-base titration, the amount of calcium carbonate in our sample would be similar to the actual amount of calcium carbonate in the antacid medication. To begin the experiment, we followed the developed standard method of acid-base titration given during the laboratory period and utilized the required materials. Following the completion of the experiment, our calculated amount of calcium carbonate was 514.93 mg/tablet, which is relatively accurate to the actual value of calcium carbonate in the antacid medication at 500 mg/tablet. Given the results, we determined our experimental methods were effective in determining the content of calcium carbonate in the antacid sample.
Diprotic Acid Abstract: In this experiment, the total amount of acid neutralized by a solution containing both Na2CO3 and NaHCO3 was determined followed by using the information to calculate the concentrations of carbonate and bicarbonate ions in the solution. This was done by equipping LoggerPro with a pH sensor and a drop counter in order to calculate the pH levels after every single drop. Three titrations were performed for analysis: a titration of a known acid with Bromothymol blue indicator, a titration of an unknown acid with Bromothymol indicator, and a titration of an unknown acid with phenolphthalein. The indicators were added for comparison between color indicators and pH titrations. The results showed that the concentration of Na2CO3 in the unknown acid is between 0.03152M and 0.03924M, and the concentration of NaHCO3 is between 0.02148M and 0.02924M.