Idea Transcript
Analysis of Vitamin C Taylor Noeller CHE142.01X 4/11/2014
Noeller |1
Introduction Vitamins are a group of small molecular compounds that are essential nutrients in many multi-cellular organisms, and humans in particular. The name "vitamin" is a contraction of “vital amine”, and came about because many of the first vitamins to be discovered were members of this class of organic compounds. And although many of the subsequently discovered vitamins were not amines, the name was retained. Fruits, vegetables, and organ meats (e.g., liver and kidney) are generally the best sources of ascorbic acid; muscle meats and most seeds do not contain significant amounts of ascorbic acid. The amount of ascorbic acid in plants varies greatly, depending on such factors as the variety, weather, and maturity. But the most significant determinant of vitamin C content in foods is how the food is stored and prepared. Since vitamin C is easily oxidized, storage and the cooking in air leads to the eventual oxidation of vitamin C by oxygen in the atmosphere. In addition, ascorbic acid’s water-solubility means that a significant amount of vitamin C present in a food can be lost by boiling it and then discarding the cooking water (Kramer et. al). A deficiency of Vitamin C leads to the disease scurvy, at one time commonly occurring during long sea voyages. British sailors combatted scurvy by carrying limes, rich in Vitamin C, on their voyages, thus engendering the name “limey.” Although the Food and Drug Administration recommends a daily intake of 60 mg of Vitamin C, Linus Pauling suggested that amounts of 1-2 grams daily are instrumental in fighting the common cold (Blake). Ascorbic acid is an important biological antioxidant (reducing agent). It helps to keep the iron in the enzyme prolyl hydroxylase in the reduced form and thereby it helps to maintain the activity of the enzyme. Prolyl oxidase, on the other hand, is essential for the synthesis of normal collagen. In scurvy, abnormal collagen is synthesized that causes skin lesions and broken blood vessels (Morasch). This experiment illustrates how titration can be used to determine the ascorbic acid content of a Vitamin C tablet containing about 500 mg of Vitamin C. The mass percentage of ascorbic acid in Vitamin C is determined by titrating the Vitamin C samples with standardized sodium hydroxide solution.
Noeller |2 In some acid-base neutralization reactions, an acid reacts with a metal hydroxide base to produce water and a salt: HX(aq)
+
acid
MOH(aq)
→
metal hydroxide
H2O(l)
+
MX(aq)
water
salt
The protons (H+) from the acid react with the hydroxide ions (OH–) from the base to form the water. The salt forms by combining the cation from the base and the anion from the acid. Acids often react with bases; the solubility of the salt does not determine whether the reaction occurs or not. When carrying out an acid-base neutralization reaction in the laboratory, you observe that most acid solutions and base solutions are colorless, and the resulting water and soluble salt solutions are also colorless. Thus, it is impossible to determine when a reaction has occurred, let alone when it is complete. If the appropriate indicator has been chosen, the endpoint of the titration (i.e., the color change) will occur when the reaction is complete, or when the acid and base are stoichiometrically equivalent (Reed): moles of acid = moles of base A Vitamin C tablet contains ascorbic acid, HC6H7O6 (aq), as well as binder material that holds the tablet together. The balanced equation for the reaction between ascorbic acid and sodium hydroxide is shown below: HC6H7O6 (aq) ascorbic acid
+
NaOH(aq) sodium hydroxide
→
H2O(l) water
+
NaC6H7O6 (aq) sodium ascorbate
The objectives in this lab were:
Determine the experimental molar mass of ascorbic acid
Determine the amount (in mg) of ascorbic acid in a vitamin C tablet
Noeller |3
Experimental Procedure
1. Placed a preweighed vitamin C tablet in a preweighed Erlenmeyer flask 2. Added 50 mL of distilled water and let tablet soften for 10 minutes 3. Used stirring rod to break up tablet (rinsed stirring rod into flask) 4. Added 3 drops of phenolphthalein and magnet 5. Titrated to light pink colored solution a. Placed a white sheet of paper under the flask b. Turned on the magnetic stirrer to stir the acid at a steady rate c. Slowly opened the stopcock and allowed NaOH to drip into flask d. Stopped the addition of NaOH when the solution turned light pink 6. Recorded values and completed calculations Data Table Data mass of flask and tablet mass of flask mass of tablet final buret reading initial buret reading mL of NaOH added L of NaOH added
90.80 g 90.20 g 0.60 g 31.20 mL 0.10 mL 31.10 mL .0311 L
Results- Objective 1 moles of NaOH added L of NaOH x M of NaOH moles of ascorbic acid neutralized stoichiometry; 1:1 mole ratio molar mass of ascorbic acid g of ascorbic acid / mol of ascorbic acid percent error (experimental – true) / true x 100%
3.11 x 10-3 mol 3.11 x 10-3 mol 192.9 g/mol 9.6%
Results- Objective 2 grams of ascorbic acid neutralized mol of ascorbic acid x g/mol of ascorbic acid milligrams of ascorbic acid neutralized g ascorbic acid x 1000 percent error (experimental – true) / true x 100%
0.547 g 547 mg 9.4%
Noeller |4
Calculations mass of tablet
90.80 g – 90.20 g = 0.60 g
mL of NaOH added
31.20 mL – 0.10 mL = 31.10 mL
L of NaOH added
31.10 mL / 1000 L/mL = 0.0311 L
moles of NaOH added
0.0311 L x 0.10 M = 3.11 x 10-3 mol
molar mass of ascorbic acid
0.60 g / 3.11 x 10-3 mol = 192.9 g/mol
percent error (obj. 1)
(192.9 g – 176 g) / 176 g x 100% = 9.6%
grams of ascorbic acid neutralized
3.11 x 10-3 mol x 176 g/mol = 0.547 g
milligrams of ascorbic acid neutralized
0.547 g x 1000 mg/g = 547 mg
percent error (obj. 2)
(547 g – 500 g) / 500 g x 100% = 9.4%
Conclusion The first objective of this lab was to determine the experimental molar mass of ascorbic acid. My experimentally determined molar mass was 192.9 g/mol. The true molar mass of ascorbic acid is 176 g/mol. Our experimental data yielded an 9.6% error. Human error that may have caused this positive percent error includes adding too little NaOH during titration. If too little NaOH was added, then my calculations would conclude a fewer number of moles neutralized than should have been concluded. With the number of moles calculated being too small, then the ratio of grams per mole would be wrong and would produce a number that is too large. This could very well be one of the reasons why my experimental molar mass is larger than the true molar mass. Another possible error is that the tablet was not pure ascorbic acid. The tablet included binders which contributed to the mass of the tablet. Given that the mass of the ascorbic acid was actually less than the mass of the whole tablet, then the ratio of grams of ascorbic acid per moles would be smaller. With a smaller mass number, the
Noeller |5 calculations for experimental molar mass could have been smaller and therefore closer to the true molar mass. The second objective was to determine the amount (in milligrams) of ascorbic acid in a vitamin C tablet. The true value for the mass of ascorbic acid is 500 mg. My experimental value was 547 mg. This gives me a 9.4% error. This positive error could be attributed to too many moles of ascorbic acid calculated. If too much NaOH was added during the titration, then my calculations would conclude a larger number of moles of ascorbic acid being neutralized. With a smaller number of moles being multiplied by the molar mass, the result would be a smaller number of milligrams that could be closer to the true value. It is also possible that the true value was wrong. The true value could have been an approximation. If the true value was an approximation, then it is possible that my experimental value was closer to correct producing a smaller percent error.
References 1. Blake, Robert. "Ascorbic Acid in Vitamin C Tablets." CHM 151L. Glendale Community College. Glendale, Arizona. http://web.gccaz.edu/~rob2108739/Ascorbic%20Acid%20Titration/Ascorbic%20A cid%20Titration%20Report%20Fall%202013.pdf 2. Kramer, B.K., V.M. Pultz, and J.M. McCormick. "Vitamin C Analysis." CHEM130. Truman State University. Kirksville, Missouri. 25 Apr 2011. http://chemlab.truman.edu/CHEM130Labs/VitaminC.asp 3. Morasch, Ralph. "Analysis of Vitamin C." Chem-131 Lab-08. Pierce College. Lakewood, Washington. http://www.pierce.ctc.edu/staff/dwoods/Chem131/Lab/Chem-131%20Lab-08%2010-1%20Analysis%20of%20Vit-C%20(std).pdf 4. Reed, Robin. "ANALYSIS OF VITAMIN C." CHEM 1021. Austin Peay State University. Clarksville, Tennessee. https://www.apsu.edu/sites/apsu.edu/files/chemistry/SP11_1021_ANALYSIS_OF_VI TAMIN_C.pdf