Wait a second!
More handpicked essays just for you.
More handpicked essays just for you.
Experiment preparation of organic compounds
Scientific method essaiy
Scientific method essaiy
Don’t take our word for it - see why 10 million students trust us with their essay needs.
Recommended: Experiment preparation of organic compounds
The goal of the project was to characterize an "unknown" organic acid in order to make a proper identification of the acid, while learning proper techniques for scientific measurement and analysis of error.
In order to ensure the most accurate data, a purification was performed by the process of recrystallization. To perform the recrystallization the powder was dissolved in a minimal amount of hot ethanol/H2O solvent that allowed the unknown powder to crystallize properly when cooled. This process allowed for the removal of soluble impurities when suction filtered. A sample of the unknown acid was weighed at 8.24 g, and it was found that 164ml of a 40% ethanol, 60% H20 solvent dissolved the 8.24 g of unknown acid when heated. The beaker containing the dissolved acid was then placed in a beaker containing ice, allowing the unknown acid to recrystallized. After vacuum filtration, the recovered unknown was dried and weighed at 6.92 g. The percent recovery was determined by the following calculation: (8.24--6.92)8.24 x 100% = 16% loss = 84% recovery of unknown.
The melting point of the unknown was determined for both pure (recrystallized) and impure (original) powders. To perform thee melting point procedure the powders were placed in small tubule. The tubule was placed in a Mel-temp device and the melting range was recorded. The data is present in the form of a range in the following chart:
"Melting Point"
Substance Starting point (°c) Ending Point (°c)
Pure Acid (recrystallized)
Pure Acid 2nd Range 125.3
7 129.2
6
Impure Acid (original) 124.1 128.2
In order in determine the equivalent weight of the unknown acid, a base was needed at a known or "stan...
... middle of paper ...
...erivative of the unknown acid, it was determined that the unknown acid 94 is sebacic acid [HOOC(CH2)8COOH] whose structure is shown below.
A comparison of the experimental value for the melting point of the melting point of sebacic acid yields an error of 5%, represented by the following calculation: (133°C-125°C)/133°C x 100% = 5%. A comparison of the experimental value for the equivalent weight of sebacic acid yields an error of 2.5%, determined by the following calculation: (101-98)/101x100% = 2.5%. A comparison of the experimental value for the melting point of the amide derivative of sebacic acid yields an error of 0%. This was determined by the following calculation (168-168)/168x100%=0. On average, the error inherent to the total project was 2.5%. In the future, it would be desirable to further support the identification of the unknown
Solid A was identified to be sodium chloride, solid B was identified to be sucrose, and Solid C was identified to be corn starch. Within the Information Chart – Mystery White Solid Lab there are results that distinguishes itself from the other 4 experimental results within each test. Such as: the high conductivity and high melting point of sodium chloride, and the iodine reaction of corn starch. Solid A is an ionic compound due to its high melting point and high electrical conductivity (7), within the Information Chart – Mystery White Solid Lab there is only one ionic compound which is sodium chloride, with the test results of Solid A, it can be concluded that is a sodium chloride. Solid B was identified as sucrose due to its low electrical
The purpose of the Unknown White Compound Lab was to identify the unknown compound by performing several experiments. Conducting a solubility test, flame test, pH paper test, ion test, pH probe test, conductivity probe test, and synthesizing the compound will accurately identified the unknown compound. In order to narrow down the possible compounds, the solubility test was used to determine that the compound was soluble in water. Next, the flame test was used to compare the unknown compound to other known compounds such as potassium chloride, sodium chloride, and calcium carbonate. The flame test concluded that the cation in the unknown compound was potassium. Following, pH paper was used to determine the compound to be neutral and slightly
Extraction is a separation method that is often used in the laboratory to separate one or more components from a mixture. Sucrose was separated at the beginning because it is the most immiscible and it’s strongly insoluble. Next Acetylsalicylic Acid was separated which left Acetanilide alone. Variety steps could have led to errors occurring. For example the step of separation, when dichloromethane layer was supposed to be drained out, it could be possible some aqueous layer was drained with it. Which could make the end result not as accurate. Also errors could have occurred if possibly some dichloromethane was not drained out. Both way could interfere with end result of figuring the amount of each component in the mixture. The solids percentage were 22.1% more than the original. That suggests that solids weren’t separated completely which clarifies the reason the melting points that were recorded were a slightly lower than the actual component’s melting point. The melting point for Acetylsalicylic Acid is 136 C but that range that was recorded during the experiment was around 105 C to 118 C. The melting points were slightly lower than the literature value. Sucrose was the purest among all component due to its higher melting point which follows the chemical rule that the higher the melting point the more pure the component
The primary goal of this laboratory project was to identify an unknown compound and determine its chemical and physical properties. First the appearance, odor, solubility, and conductivity of the compound were observed and measured so that they could be compared to those of known compounds. Then the cation present in the compound was identified using the flame test. The identity of the anion present in the compound was deduced through a series of chemical tests (Cooper, 2009).
Purpose: To identify the mystery powder based on its physical and chemical properties, comparing them to the five substances and which one matches. The chemical reactions with water, universal indicator, vinegar and Iodine solution are then analyzed and matched with each other to determine the mystery powder.
Since, the expected weight was 50.63 mg the percent yield is 59.3%. A TLC was conducted on this final product and a faint spot of 4-tert-butylcyclohexanone still appeared in lane 3 of the plate; meaning the reaction did not fully go to completion. The Rf values were 0.444, 0.156, and 0.111, where the lowest value is the trans isomer and the highest value is the ketone. This affected the IR spectrum conducted by having a carbonyl group peak at 1715 cm-1 which should not be present if all the product was 4-tert-butylcyclohexanol. However, the IR spectrum still showed peaks at 3292 cm-1 (hydroxyl group), 2939 cm-1 (sp2 carbon bonded to hydrogen) and 2859 cm-1 (sp3 carbon bonded to hydrogen) which support the presence of the alcohol. The accepted melting point of 4-tert-butylcyclohexanol is in the range of 62 – 70˙C (Lab Manual). The two melting point measurements using the Mel-Temp® machine gave ranges of 57 – 61˙C and 58 – 62˙C, which is not exact due to some 4-tert-butylcyclohexanone being present that has a low melting point of around 47 – 50˙C
Mixed melting point was used to confirm the identity of the product. The smaller the range, the more pure the substance. When the two substances are mixed; the melting point should be the same melting range as the as the melting range obtained after filtering. If the mixed melting point is lower one taken from the crystals, then the two substances are different.
Also, looking at Table 1, the percent yield is shown to be 44.9%. The percent yield is how much product was recovered after the reaction was carried out. The percent yield can be used to explain why the melting point observed in the experiment didn’t match the known melting point. Obtained melting points are generally lower than the literature value melting points of a substance due to the number of impurities present in the obtained product. The percent yield of 44.9% validates that the product could have had some impurities present, and thus the lower melting point.
Overall this experiment was a success yielding 98.8% of the initial 1.34g of known compounds. Looking at Table 1 the problem of separation quickly becomes apparent, both M-Toluic Acid and Acetanilide are insoluble in water. This left two non-salts in one mixture, and what solvent to use to separate these two was the most important question as their respective melting points are also very similar. After looking at both compounds and noticing the M-Toluic Acid (Image 2) had an OH group hanging off of it next to a double bond, the H ion on the end would be susceptible to a base. But further investigation showed the large number of hydrogen atoms hanging off the Acetanilide (Image 3) and it was thought that the NaHCO3 would be strong enough to rip the Acetanilide apart.
an unknown amino acid. A titration curve is the plot of the pH versus the volume
Aspirin has a white colour, a crystalline structure and is weakly acidic. The melting and boiling points are 136 °C and 140 °C respectively.
After synthesizing a chemical, especially a drug, it is important to confirm the identity and purity of the product. You will perform three tests to examine the identity and purity of the aspirin that you synthesized. One test will detect the presence of leftover salicylic acid in the synthesized aspirin and allow you to determine its concentration. Government regulations stipulate that commercial aspirin must not contain residual salicylic acid since it is irritating to the mouth, throat, and stomach. Iron salts react with phenols to form a complex ion that has a purple color, therefore iron (III) chloride can be used to determine if your aspirin sample contains residual salicylic acid. The second test uses melting point to evaluate the purity of your aspirin product. You will measure the melting point of pure acetylsalicylic acid (135°C) as a comparison to your product. The melting point of a pure aspirin sample should be within 1°C of its known melting point. A compound that contains impurities will tend to melt over a range of temperatures and at temperatures lower than the fixed mp for the pure compound. For example naphthalene, an ingredient of mothballs, has a melting point of 80°C and a pure sample of naphthalene would most likely be observed to melt within 80-81°C. An impure sample of naphthalene might melt over a range from 75-80 °C. Impurities in the crystals of the compound weaken the structure, which results in the compound melting at a lower temperature. In addition, impurities will be unevenly distributed throughout a sample of a compound. This non-uniform composition results in some areas of the sample being more “pure” than othe...
The conical vial was placed in a small beaker and allowed to cool to room temperature. The mixture was Cooled thoroughly in an ice bath for 15-20 minutes and crystals collected by vacuum filtration on a Hirsch funnel. The vial was rinsed with about 5 mL of ice water and transferred into to the Hirsch funnel and again washed with two additional 5mL portions of ice water. Crystals were dried for 5-10 minutes by allowing air to be drawn through them while they remained on the Hirsch funnel. The product was transferred to a watch glass plate and allow the crystals to dry in air. Crude acetaminophen product was weighed and set aside a small sample for a melting point determination and a color comparison after the next step. Calculation of the percentage yield of crude acetaminophen (MW = 151.2). was done and recorded in the lab notebook.
Based on the results obtained from the experiment, the equivalence volume of the titration of the weak acid (HC2H3O2) and strong base NaOH was 28.11 mL, the calculated Ka from the pH of the original solution was 6.92 x 10-6, and the calculated Ka from the pH of the Half-Neutralization point was 1.23x10-5. There was a 60.5% error in the calculated Ka from the pH of the original solution and 29.7% error in the calculated Ka from the pH of the Half-Neutralization point, indicating that there could have been numerous sources of
After the some time, we filter it through a Büchner funnel before it is recrystallized and filtered again. The mass was recorded as it was dry. By adding sodium carbonate, we will now test whether what obtained is benzoic acid or not, because one can observe bobbles if it is an acid. After that we burn it to test if it is aromatic.