The ascorbic acid, or vitamin C (C6H8O6), content

RKPAdapted from UCSD6CL Experiment 6rev 1/01SOUTHWESTERN COLLEGEChemistry 250 – Analytical ChemistryThe Determination of Ascorbic Acid in Vitamin TabletsINTRODUCTIONThe ascorbic acid, or vitamin C (C6H8O6), content of vitamin tablets can be determined by avariety of techniques. This experiment will utilize a series of redox steps for analysis. Veryinteresting oxidation/reduction chemistry can be studied along the way. A back titration will benecessary for the final step.This experiment will require three standard solutions: KIO3, KBrO3, and Na2S2O3. The first twocan be obtained as solids in pure forms that can be easily dried (without decomposition) andweighed and therefore make excellent primary standards. Sodium thiosulfate, however, isusually obtained in the pentahydrate form, Na2S2O3 .H2O. The crystalline lumps are usually atleast partially opaque due to significant loss of the waters of hydration, which increases (but notquantitatively) with heating. Therefore, since the exact composition is not known, it cannot beused as a primary standard and must be titrated to determine the concentration.PART IIn the first part of the experiment, a standardized sodium thiosulfate solution is prepared usingKIO3 as the primary standard. A standard amount of KIO3 is added to excess KI to generate thetri-iodide ion according to the reactionIO3- +8 I-+6 H+3 I3-+3 H2O(1)The I3- (often referred to as iodine in solution even though elemental I2 is only very slightly solublein water) is the major species when I2 is in an aqueous solution of I-.I2+I-I3-K = 270(2)The triiodide ion is a weak (and therefore selective) oxidizing agentI3-+2 e-3I-Eο = +0.54 V(3)and can be used without an indicator if the concentration is high enough since the triiodide ion isdeep red-brown at high concentrations and the iodide ion is colorless. In dilute solutions,however, the transition color becomes pale yellow to clear and hard to detect. Starch and the2triiodide ion make a deep blue colored complex, which can be used to enhance the endpointdetection (now the deep blue to colorless). BEWARE: starch decomposes quickly and MUST beprepared fresh frequently.The amount of triiodide ion formed from reaction (1) is determinable if a known mass (andtherefore known number of moles) of the primary standard KIO3 is reacted with excess I- since thisreaction goes quantitatively. The I3- can then be used to standardize a thiosulfate solution by thefollowing reaction:I3- +S4O62-2 S2O32-+3I-Eο = +0.46 V(4)This reaction is run in the presence of starch. Since both the thiosulfate and the tetrathionate(S4O62-) ions are colorless, the endpoint can be detected as the blue to colorless change due to thedisappearance of the starch-triiodide complex. The I3- starch complex is very strong and thereforeis slow to dissociate as the stoichiometric endpoint nears. Therefore time must be allowed for itsdissociation. The E° value is large enough to ensure quantitative results. NOTE: the blue colormay reappear with time due to the air oxidation of the iodide ion.Solutions of Na2S2O3 are prepared from the solid and include a small amount of Na2CO3 whichraises the solution pH to improve the stability and to precipitate out any trace amounts ofcopper(II) that might be present, since this ion acts as a catalyst in the decomposition of thiosulfateion. It is also a good idea to store the solution in the dark to slow its decomposition. Unlike moststandard acids and bases, however, solutions of Na2S2O3 have a limited shelf life.NOTE: The stoichiometry for the calculations must involve a combination of that from (1) and(4) with KIO3 as the limiting reagent.PART IINow that the Na2S2O3 has been standardized, it can be used indirectly to determine the vitamin Ccontent of vitamin tablets. First the tablets must be crushed and the mass determined. Oftenbinders are present and remain suspended but do not affect the results. In some tablets, the bindermay be starch so that the characteristic color of the complex with I3- may be seen early in theanalysis, but this will not affect the results.An excess of Br- is added to the solution containing vitamin C and a measured amount of standardKBrO3 solution is added. The product, Br2, reacts with vitamin C to oxidize it.BrO3- +5 Br-Br2(Vit C)red++6 H+3 Br2(Vit C)ox ++2 Br-3 H2O+2 H+(5)(6)An excess of BrO3- is added so that all of the vitamin C reacts and a quantitative excess of Br2remains. This amount must now be determined.3Excess KI is added to this to form the triiodide ionBr2+3 I-2 Br-+I3-(7)which is titrated with the thiosulfate standard from Part I above.I3-+2 Na2S2 O3S4O62- +3 I- +4 Na+(8)The endpoint is visually seen as the disappearance of the blue I3- starch complex color.This experimental procedure assumes that the unknown samples will contain about 100 mg ofvitamin C. This is the approximate contents of 1 cup (250 mL) of orange juice, 1 cup of broccoli(extracted), one multiple vitamin (beware of other species in a multiple vitamin mix that mightinterfere with the analysis), or one 100 mg vitamin C tablet. The current RDA (recommendeddaily allowance) value for vitamin C is 60 mg.The calculations are somewhat involved, so an outline for the procedure is given at the end of theexperiment. To make certain that you understand these, try the prelab questions before beginningthe experimental work.PRELABORATORY QUESTIONS1. If in Part I Step 2, you measured 0.2287 g of KIO3, what would be the molarity of the solution?2. If in Step 4, 18.21 mL of the sodium thiosulfate solution were used for the titration, what is themolarity of the standardized thiosulfate solution?3. If in Part II Step 1, you measured 0.1156 g of KBrO3, what would be the molarity of the solution?4. If Step 5 took 1.82 mL to back titrate, how many mg of vitamin C were in the sample?5. Calculate the equilibrium constant for equation 4. (HINT: you know Ε°, and from the lecture youknow the relationship between Ε°, ΔG, and K.)*** Check your answers with your instructor before proceeding ***4EXPERIMENTAL PROCEDURESafety Precautions♦ Wear safety glasses at all times when in the laboratory.♦ Exercise caution when dealing with strong acids and bases as they are extremelycorrosive and poisonous. Avoid contact with your skin, eyes, or clothes. If youspill the any on yourself, rinse thoroughly with water and contact your instructorimmediately.♦ Solid KBrO3 can cause a fire if it comes in contact with damp organic materials,like paper towels.PART I1.Prepare a solution of approximately 0.10 M Na2S2O3. Boil 150 mL of water for 15 minutes.Allow it to cool to room temperature. Then add 3.8 g of Na2S2O3.5H2O and 0.20 g of Na2CO3.Stir well. Transfer to a plastic or glass bottle and store in the dark. Do not let excess light reachthe solution. It is best to let it sit a day before using.2.The stockroom will pre-dry (1hour at 110°C some KIO3 for you. DO NOT contaminate this. Donot put your spatula in the bottle or return any unused solid to the bottle. Make certain to leave thelid on as much as possible so it does not absorb water. Measure about 0.22 g (to the nearest 0.1mg) of the cooled salt and transfer quantitatively to your 100 mL volumetric flask. Dissolve anddilute to the mark. Calculate the molarity of this (it should be close to 0.01 M).3.The starch solution was prepared for you by dissolving 1 gram of soluble starch in 15 mL of waterto make a paste. It was diluted to 500 mL with boiling water and heated until clear. A pinch ofHgI2 was added to keep bugs from growing. Please keep the bottle tightly stoppered.4.Standardize your sodium thiosulfate solution. Pipet 25 mL of the standard iodate solution into a125 mL Erlenmeyer flask. Add 1 gram of iodate-free KI and swirl to dissolve. Add 1 mL of 6 MHCl and titrate immediately with the thiosulfate solution until only a pale yellow exists (thisoccurs just prior to the stoichiometric point). Add 5 mL of starch solution (it must turn blue nowor you added too much thiosulfate) and titrate slowly until the blue color just disappears.5.Repeat the standardization twice more (for a total of three). Process each sample totally beforebeginning the next one to minimize air oxidation of the iodide ion.6.Calculate the concentration of the Na2S2O3 solution (watch the stoichiometry in your calculations).Check with your instructor if the three values are not close to each other.7.Dispose of all solutions in the proper containers in the hood.5PART II1.The stockroom will supply predried KBrO3 for you (dried at least one hour at 110°C, DO NOTCONTAMINATE). Measure approximately 0.15 g (to the nearest 0.1 mg) of the cooled salt andtransfer quantitatively to a 100 mL volumetric flask. Dissolve and dilute to the mark. Calculatethe molarity of this (it should be about 0.009 M). WARNING: solid KBrO3 can cause a fire if itcomes in contact with damp organic materials, like paper towels. Dispose of any excess correctly!See your instructor.2.Prepare the sample for ascorbic acid analysis. If a vitamin tablet is used, pulverize it with a mortarand pestle and transfer quantitatively to a volumetric flask, which will hold approximately 1 mLfor each mg of vitamin C in the tablet. Add a spatula-tip full of solid EDTA disodium salt to thecontents of the flask. This will tie up trace Cu2+ that can catalyze vitamin C decomposition.Dilute to the mark with 1.5 M H2SO4. Use 25 mL aliquots of this for your trials.3.Do the following process completely for each sample before proceeding to the next one.4.If a diluted tablet was used, proceed with a 25 mL aliquot of the sample. If a solid sample wasused, dissolve the sample in 25 mL of 1.5 M H2SO4. If a liquid juice was used, add 25 mL of1.5 M H2SO4 to 50.0 mL of juice. In any case, proceed as follows. Add 2.5 g KBr (this amount isexcess and need not be exact) and immediately add 25.00 mL of the standard KBrO3. A faint (butdefinite) yellow color must be seen (due to excess Br2). If not, add more (a known amount, ofcourse) of KBrO3 until yellow.5.Add 1.5 g of KI (not critical and need not be exact) and 5 mL of starch and back titrate with thestandardized thiosulfate solution. Be careful. This volume should not be over a few milliliters.6.Repeat steps 4-5 with your second sample and then your third.7.Calculate the mg of ascorbic acid per tablet or per mL, of juice.CALCULATION OUTLINE FOR PART IIA.Calculate the concentration of the primary standard solution of KBrO3 (you know the mass, molarmass, and volume to which it was diluted).B.Calculate the back titration.a.b.C.Knowing the volume of Na2S2O3, and its concentration from Part I, calculate the moles ofthiosulfate for the back titration.Using stoichiometry (equations 8, 7, and 5) equate this to the number of moles of excessBr2 in the solution.Calculate the moles of BrO3- added (you know the volume and calculated the molarity in Aabove), and from this the moles of Br2 formed in the solution.6D.Subtract B from C to get moles of Br2 actually used for the vitamin C.E.Using stoichiometry (equation 6) convert the answer in D to moles of vitamin C present.F.Knowing the molar mass of vitamin C, convert the answer in E to mass of vitamin C.G.Depending on your procedure for processing your original tablet, now calculate the mass andpercent of the RDA value of vitamin C for each full tablet, or for each 250 mL glass of juice.POSTLABORATORY QUESTIONS1.Often vitamin C is determined by direct titration with a standardized solution of 12 (in excess Ithat produces I3-). Balance the equation between vitamin C and the triiodide ion.2.If 21.35 mL of a 0.0215 M I2 solution (in excess KI) is needed to reach the endpoint (see the aboveproblem) when a 25 mL aliquot of 100 mL of solution made from a vitamin tablet, how many mgof vitamin C are in the tablet?

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