Orgo Post-Lab Questions and Overview

(A) Calculate the volume of 6M NaOH required to react completely with the amount of methyl salicylate you used. How much of the 6M NaOH that you used was in excess of the theoretical amount? (B) What volume of 3M H2SO4 is needed to neutralize all of the disodium salicylate and the excess NaOH present after the initial reaction? How much sulfuric acid was in excess?

(A) 1.498 g Methyl Salicylate x (1 mol/ 152.1 g) x (2 mol NaOH/ 1 mol) x (1000 mL/ 6 mol) = 3.283 mL NaOH 15 mL NaOH used - 3.283 mL needed = 11.717 mL excess NaOH (B) 1.498 g Methyl Salicylate x (1 mol/ 152.1 g) x (1 mol Disodium Salicylate/ 1 mol) x (1 mol H2SO4/ 1 mol) x (1000 mL/ 3 mol) = 3.283 mL H2SO4 needed to neutralize disodium salicylate 11.717 mL NaOH x (6 mol/ 1000 mL) x (1 mol H2SO4/ 2 mol NaOH) x (3 mol/ 1000 mL) = 11.717 mL H2SO4 needed to neutralize NaOH 3.283 + 11.717 = 15 mL needed total... 16 mL used - 15 mL needed = 1 mL excess

Describe and explain the possible effect on your results of the following experimental errors or variations. (a) You added only enough 3M H2SO4 to bring the pH down to 4. (b) Your reaction mixture had an oily layer on top when you added the 3M sulfuric acid. (C) Because of a label-reading error by a lab assistant, the bottle labeled "salicylic acid from benzene" actually contained acetylsalicylic acid (aspirin)

(A) If the pH was not reduced to at or below 2, H2SO4 may become the limiting reagent and the concentration of H^+ in solution may be too low to drive the reaction. This would result in a very slow reaction with a decreased product yield. (B) This oily layer on top is likely unreacted reagent still present in solution. This indicates that heat of reflux was not carried out long enough. This layer may have prevented the sulfuric acid from reacting with the disodium salicylate, also decreasing the product yield. (C) Due to melting point depression, adding acetylsalicylic acid instead of salicylic acid wouldve lowered and broadened the melting point significantly and any conclusions based off of this data would be incorrect.

(a) What is the minimum volume of boiling water needed to dissolve 0.200g of phenacetin? (b) About how much phenacetin will remain dissolved when water is called to room temperature? (c) calculate the maximum mass of solid phenacetin that can be recovered when the cooled solution is filtered.

(a) 0.200 g x 100 mL / 1.22 g = 16.4 mL of water (b) 16.4 mL x .076g / 100mL = .0125 g of phenacetin (c) 0.200 g - 0.0125 g = .1875 g

Write balanced reaction equations for the reactions involved (a) when aspirin dissolves in aqueous NaHCO3, and (b) when aspirin is precipitated form a sodium acetylsalicylate solution by HCl. Assuming that both reactions are spontaneous under standard conditions, label the stronger acid, stronger base, weaker acid, and weaker base in each equation.

(a) C9H8O4 (s) + Na(+)HCO3(-) → Na(+)C9H7O4(-) + H2CO3. (b) C9H7NaO4 + HCl → C9H8O4 + NaCl + H20 (a) HC9H7O4(s) + NaHCO3(aq) => NaC9H7O4(aq) + H2CO3(aq) stronger acid stronger base weaker base weaker acid pKa HC9H7O4 (3.49) < pKa H2CO3 (6.35), so HC9H7O4 is the stronger acid compared to H2CO3 and NaC9H7O4 (conjugate base of HC9H7O4) is the weaker base compared to NaHCO3 (conjugate base of H2CO3). (b) NaC9H7O4(aq) + HCl(aq) => HC9H7O4(s) + H2O(l) stronger base stronger acid weaker acid weaker base HCl is a strong acid (fully ionized) while HC9H7O4 is a weak acid (partially ionized), so HCl is the stronger acid compared to HC9H7O4 and H2O (conjugate base of HCl) is the weaker base compared to NaC9H7O4 (conjugate base of HC9H7O4).

Describe and explain how each of the following experimental errors or variations might affect your results. (a) You failed to dry the reaction flask after washing it with water. (b) You forgot to add the sulfuric acid. (c) You used twice the amount of acetic acid specified in the procedure (d) You left out the sodium bicarbonate washing step (e) Your thermometer bulb was 1 cm higher than it should have been

(a) If one had failed to dry the reaction flask after washing it with water, there would be water present in the initial reaction of isopentyl alcohol and acetic acid. This water would increase one product of the reactant which would make the equilibrium favor the reactants. This would lower the recovery of isopentyl acetate. (b) If the sulfuric acid had not been added, the reaction would have been without its catalyst. Thus, leaving the reaction much slower. The recovery of the reaction would be significantly decreased. (c) If twice the amount of acetic acid was used, following Le Chatlier's Principle, the equilibrium would have been driven to produce more products. This would increase the recovery of isopentyl acetate. (d) If the sodium bicarbonate washing step had been skipped, there would be unwanted acid and byproducts present in the reaction. Water removes the acid and then the sodium bicarbonate finishes the job. This would falsely increase the recovery. (e) If the thermometer were too high, the temperature would not read accurately the temperature of the condensing vapors during the boiling point range. This would have lowered the observed boiling point range, making the drops collected at each temperature inaccurate. The first ten drops would still be theoretically isopentyl alcohol, and the next drops (between 137-142) would still be the product. However, because the temperature would read artificially low, the drops after the actual temperature of 142 would also be collected as product when in reality they should be discarded. This would artificially increase the recovery.

Tell how each of the following experimental errors will affect your experimental results (yield, purity, or both), and explain why. (a) You failed to dry the product completely. (b) You used enough water to recrystallize phenacetin, but your unknown was acetanilide. (c) In the "Extraction and Evaporation" Experiment, you didn't extract all of the aspirin from the DCM solution.

(a) If the product were not completely dried the yield would artificially go up because the presence of water would increase the mass, and the purity would go down as there is water and unknown in your product. (b) Phenacetin requires a large amount of water while acetanilide requires much less. Adding the amount of water for phenacetin would significantly decrease your yield of acetanilide and decrease your purity. (c) Because aspirin does not dissolve in DCM, if you failed to extract all from the DCM, when evaporating the DCM the solid would contain both the unknown and aspirin. This would increase your yield of unknown while decreasing the yield of aspirin and decreasing your purity.

Tell whether each of the experimental errors in Exercise 3 will affect the melting point of the unknown component. If it will, tell how it will affect the melting point, and explain why.

(a) The melting point would be lowered and broadened because water makes the solid impure and has a very low melting point. (b) The melting point would not be effected. (c) Aspirin has a melting point of 136 degrees Celsius, which is comparable to that of the unknown (acetanilide). This being said, as an impurity, this would lower and broaden the melting point of the acetanilide.

(a) Describe any evidence that a chemical reaction occurred when you added 6M HCl to the solution of sodium acetylsalicylate. (b) Explain why the changes that you observed took place.

(a) When 6 M HCl was added to the sodium acetylsalicylate solution a white and cloudy precipitate formed. (b) After addition of the HCl, the sodium ion on sodium acetylsalicylate left to bond with the chlorine anions in solution. The hydrogen from the hydrochloric acid attached to the aspirin anion to satisfy the oxygen. The end product was the solid aspirin, which precipitated, leaving aqueous sodium chloride.

Based on the amount of water you used during the recrystallization, estimate the amount of salicylic acid that was lost as a result of being dissolved in the filtrate. Assume that the recrystallization solution was called to 10 C; the solubility of salicylic acid at that temperature is 0.14 g per 100 mL water.

0.14 g / 100 mL H2O = X / 30 mL H2O , X= .042 g Salicylic Acid

Describe and explain the possible effect on your results of the following experimental errors or variations. In each case, specify the component(s) whose percentage(s) would be too high or too low. (a) After adding DCM to Panacetin, you didn't stir or shake the mixture long enough. (b) During the NaOH extraction, you failed to mix the aqueous and organic layers thoroughly. (c) You mistakenly extracted the DCM solution with 5% HCL rather than 5% NaOH. (d) Instead of using pH paper, you neutralized the NaHCO3 solution to pH 7 using litmus paper.

2. (a) The soluble components of Panacetin would not have been dissolved completely, therefore the percentage of sucrose recovery would be artificially increased and the recovery of aspirin and acetanilde would be artificially lower. (b) Failure to mix the aqueous and organic layers would result in the inability of the organic layer to mix with the aqueous layer to convert the aspirin to its salt form for extraction. This would make the percentage of sodium acetylsalicylate (aspirin upon addition of H+) recovery too low while the percent of acetanilde too high. (c) Because HCl could not deprotonate the aspirin in the DCM solution, the aspirin salt would not have been extracted, leaving the percent recovery of aspirin at zero and the percent recovery of acetanilde artificially increased. (d) Only reducing the pH to 7, instead of 2, may not allow all aspirin to precipitate out and be recovered. This would lower the percent recovery of aspirin.

What gas escaped during the sodium bicarbonate washing? Write balanced equations for two reactions that took place during this operation.

Carbon dioxide (CO2) escaped during the sodium bicarbonate washing. (1) NaHCO3 (s) + CH3COOH (l) → CO2 (g) + H2O (l) + Na+(aq) + CH3COO^- (aq) (2) H2SO4 (aq) + NaHCO3 (aq) → HSO4^-(aq) + H2O (l) + CO2 (g) + Na+(aq)

Experiment 4 purpose

Carry out an organic synthesis, the preparation of salicylic acid from methyl salicylate.

The water solubilities of oxalic acid and sodium oxalate at room temperature are 10g / 100 mL and 3.7 g / 100 mL, respectively. Could you prepare oxalic acid by adding HCl to a solution of sodium oxalate, cooling it to room temperature, and filtering the resulting mixture? Explain why or why not.

Heating a substance is one way of increasing its solubility. When hydrochloric acid is added the heated sodium oxalate solution it reacts to form oxalic acid. The solution is then cooled as to reduce the amount of oxalic acid dissolved, increasing the amount of solid after filtration. Ideally, this solution should be cooled to near freezing to improve the percent yield. Essentially yes, this method is possible, but it will not be as efficient or accurate as if it were cooled to near freezing.

Experiment 5 Tests involved

Separatory funnel, distillation, gas chromatogprahy, IR, HNMR

Reaction scheme for experiment 1

Sodium benzoate + Hydrochloric Acid → Benzoic acid + Sodium chloride

Experiment 4 Conclusion

The experiment was successful in discovering that both derivations of salicylic acid share the same physical properties and therefore do not differ from one another.

Conclusion of Experiment 2&3

The percent compositions were calculated to be 38.23% sucrose, 27.58% aspirin, and 34.19% unknown. These do not match the theoretical values given on the label, and suggest it is inaccurate.

You can confirm a substance has been converted to a different substance by showing that the substances have different properties. Your observations should have revealed a difference in at least one property of sodium benzoate and benzoic acid. What property was that, and what observation revealed the difference?

The property difference of sodium benzoate and benzoic acid lies in their water solubilities. Sodium benzoate has a water solubility of 61.2 g/mL while benzoic acid has a water solubility of 0.34 g/mL. Sodium benzoate will dissolve greatly at room temperature while benzoic acid will dissolve very little.

Conclusion of Experiment 1

The purpose of part A of the lab, to simulate stomach acid and determine if benzoic acid is formed in the stomach, was completed. Through manipulation of the solubilities of the reagents as well as vacuum filtration, it was determined that the stomach acid pH of ~2 is low enough to cause the conversion of sodium benzoate to benzoic acid. This is proven in the end- product precipitation of benzoic acid at this pH. The cooling process allowed the benzoic acid to precipitate and be separated out, demonstrating that it was produced and present in the reaction mixture. The percent yield from this experiment was found to be 81.3%, and the percent error was 18.71%. This could result from some of the benzoic acid remaining dissolved in the reaction mixture and would not been recovered or weighed. This would falsely lower the percent yield. Also in part A, if there were an error or inaccuracy with the thermometer used and the reaction had not been cooled to near 10 Celsius, then the benzoic acid would have precipitated less. This would also result in a falsely low product of benzoic acid

Label the stronger acid, stronger base, weaker acid, and weaker base in Equation 1 for the proton transfer reaction in part A of the experiment.

The stronger acid is hydrochloric acid while the weaker acid is benzoic acid. The stronger base is sodium benzoate while the weaker is sodium chloride. This follows the concept that a stronger acid transfers protons to a stronger base to yield a weaker acid and a weaker base.

Experiment 5 Conclusion

This enabled the first drops of isopentyl alcohol, which distilled first due to its boiling point and lack of hydrogen bonding, to be distilled and not collected in the final product. Nothing after the observed boiling point was recorded in case there were unknown higher molecular weight impurities. The distillate was clear and did not contain any visible water droplets, therefore it was not dried again with sodium sulfate. It smelled of bananas, strongly suggesting the formation of isopentyl acetate. The distillate was massed for calculations, and a portion was used to perform a GC, HNMR, and IR test.

Tests employed in Experiment 2&3

Through manipulation of the solubilities of aspirin and sucrose in dichloromethane (DCM), gravity filtration of sucrose, separation of aspirin using NaOH to produce aspirin's salt (sodium acetylsalicylate) form from DCM in a separatory funnel, vacuum filtration of aspirin (as sodium acetylsalicylate) after addition of hydrochloric acid (seen in Reaction 1), and evaporation of unknown in the DCM solution, the unknown was separated from the sucrose and aspirin constituents.

Purpose of Experiment 2&3

To determine if the composition of Panacetin as stated in the label is accurate & To purify, identify, and find the final purity of the unknown substance

Purpose of Experiment 1

To place sodium benzoate into a medium that simulates stomach acid to see whether or not a new substance forms in that environment.

Experiment 5 purpose

To synthesize isopentyl acetate by combining isopentyl alcohol with acetic acid and sulfuric acid and then heating the reaction mixture under reflux for an hour

An unknown compound X is one of the four compounds listed in Table 2. A mixture of X with benzoic acid melts at 89 C, a mixture of X with phenyl succinate melts at 120 C, and a mixture of X with m-aminophenol melts at 102 C. Give the identity of X and explain your reasoning.

X is phenyl succinate because in mixture with phenyl succinate, the boiling point of phenyl succinate, 120 degrees, is preserved. (There is no boiling point elevation)

(a) Calculate the ratio of dissolved benzoic acid to benzoate ion that will exist in solution at equilibrium at a pH of 2.00. The acid equilibrium constant (Ka) for benzoic acid is 6.46 x 10 ^-5. Note that this calculation doesn't account for the benzoic acid that has precipitated from solution, but you can assume that the higher the ratio, the more benzoic acid will precipitate. (b) Carry out the same calculation for a pH of 4.00, and explain why it was so important to reduce the pH to below 4 in this experiment.

a. [H+] / Ka = [benzoic acid] / [benzoic ion], [10^-2] / [6.46 E-5], = 155 b. [H+] / Ka = [benzoic acid] / [benzoic ion], [10^-4] / [6.46 E-5], = 1.55 It is important to reduce the pH to below 4 because at the lower pH the amount of benzoic acid is higher than benzoate ion. Hence, the buffer can only be made with the use of benzoic acid.

Tests employed Experiment 4

heat reflux, vacuum filtration, Mel-temp

Tests employed in Experiment 1

vacuum filtration