3-Heptanone Lab Report

When doing the experiment, some of the mass would be loss at some points. We can improved the percent recoveries by using more solvents to remove the crystals from the flask as much as possible. Its possible that small substances of the crystals were still in the flask. 3. Were the recrystallization effective (did purification result)?

The theoretically yield for this experiment, based off of the mass of aluminum pieces originally weighed out, was 19.0144 g, while our actual yield was 12.7222 g. This is a 66.91% yield. Ideally, of course is a 100% yield. Factors that may have caused this in our own experiment include, but may not be limited to: 1) There was barely any hydrogen gas being produced in the first step when we took it off the heat, but enough to infer that the reaction may not have been completely over, possibly affecting the amount of potassium aluminum hydroxide produced 2) Having to filter the reaction mixture in the first step more than once, because we neglected to turn on the vacuum the first time, we may have lost some of the mixture in the process. 3)

There are however some aspects of the experiment that could be changed with little to no effect on the color of the solution or the calculated molar masses. For instance, if the amount of deionized water used were to be doubled, there would be no change in the results of the titration calculations. This is the reason that the amount of DI water used in this experiment was not measured nor

There are many convenient uses of TLC including monitoring the progress of a reaction, identifying a product by establishing that two compounds are the same, and determining the number of compounds in a mixture. In this experiment, TLC analysis helped determine whether the product matched any of the possible known products. One measure that TLC does not account for is the calculated percent yield. While TLC can help indicate a compound’s purity, it is not an absolute quantitative method. Only qualitative data is provided by TLC.

Introduction In recent years, there is an increasing demand of flavoured compounds in industry sectors, especially food and beverage, cosmetic and pharmaceutical industries 1–4. These compounds are generally short chain ethyl esters which are characterized by their strong fruity flavour and fragrance. Ethyl hexanoate is such a short chain acid ester which gives an apple-pineapple flavour 5,6. Most of the flavours are extracted from their natural source, but this process can take a long time and rigorous efforts eventually ending up with an inefficient yield.

The experiment that was done was to figure out “Does the amount of calcium chloride affect the temperature of water?” For the procedure, the experiment asked to record the initial temperature of 75 mL of water. The first trial said to add zero scoops of calcium chloride and stir for two minutes to record the temperature. Once the first temperature was recorded, it must be written from the difference between the initial temperature and the new temperature. Next, it asked to add one scoop of calcium chloride and stir for two minutes and record.

As seen in table 1, the theoretical yield was .712 g of C_17 H_19 NO_3. The % yield of this experiment was 7.51 % of C_17 H_19 NO_3. . This low yield can be explained from a poor recrystallization technique combined with potential contamination. Throughout the experiment, the mixture changed color from green, orange, to yellowish lime, and eventually clear.

Nevertheless, the latter is not used in this experiment since it is very reactive and extremely flammable. On the contrary, NaBH4 is relatively mild and it can be used with protic solvents. In this manner, 1.507 grs of the ketone 9-fluorenone were mixed with 30.0 ml of 95% ethanol in a 125 ml Erlenmeyer flask. The bright yellow mixture was stirred during 7 minutes until all the components were dissolved.

After all of the time data was collected, the molarity of each amount had to be recalculated, (refer to Table 3 for new molarities). Molarity (m/L) Acetone HCL Iodine Reaction 1 0.6666667 0.1666667 0.0011111 Reaction 2 1.3333333 0.1666667 0.0011111 Reaction 3 2 0.1666667 0.0011111 Reaction 4 0.6666667 0.3333333 0.0011111 Reaction 5 0.6666667 0.5 0.0011111 Reaction 6 0.6666667 0.1666667 0.0022222 Reaction 7 0.6666667 0.1666667 0.0033333 Table 3: Corrected molarity for acetone, hydrochloric acid and iodine for each reaction Next, the iodine concentration rate for each reaction needed to be calculated.

(150.22g/mol)(3.5 x 10^-3 mol of nucleophile) = 0.525 g Actual yield = 0.441 g, Percent Yield = (0.441g/0.525g) x 100% = 84% 10. Percent recovery from recrystallization = (0.172g/0.441g) x 100% = 38% 11.

The precipitate, which gets separated, was filtered, dried and recrystallised from 95% hot ethanol. Table 1: Quantity of aldehydes taken S. No. Aldehydes Molecular weight Quantity taken

The final product weight for percent yield was only the solid E product, which missed one half of the final product produce. If both products were weight, the percent yield would have been larger that it was. Instead of 22.33%, it could have been 44.66%. To prove that both products were obtained, but only one of the two products was analyze, a TLC plate of the DCM layer, that contains both products, and of the final product, was obtain.

The % yield was greater than 100% because the actual yield was greater than the theoretical yield. One error that may have caused this result to occur was that the copper may have not completely dried over the time it was left. If water was still inside the copper, it would increase the mass of the actual yield. If the copper was left in the beaker to dry for a longer time, it would help decrease the mass of the copper (as the water would be completely dried out) and bring the actual yield down.

Your actual yield is the measurement you get using a balance. Baking soda plays a huge role in this lab. Baking soda’s chemical name is sodium bicarbonate, and its chemical formula is NaHCO3. In addition, it is used in making dough rise, making soda fizzy, and is inside of our toothpastes. So when we

The moles of NaOH was then divided by the liters of NaOH used to find the molarity, which was added to trial 2 then divided by 2 to find the average