chemistry report and need the explanation and answer to help me learn.
complete the prelab and report form(“Analysis of Vitamin C via Reaction with Iodine—A Redox Titration”)r
report file, prelab and explanation of whole procedure is attached in a single file.(1st file)
Requirements: 2-4 pages
Exp 6. Analysis of Vitamin C via Reaction with Iodine—A Redox Titration CHM 121, Oakton College (last modification, Fall 2022) 1 Summary Students will standardize a solution of aqueous iodine using samples of pure ascorbic acid (Vitamin C), and then use this standardized solution to determine the amount of Vitamin C in a beverage(s) of their choice. Introduction Titrations—Applied Stoichiometry [See last week’s experimental write-up!] Molarity—A Convenient Unit of Concentration for Chemists [See last week’s experimental write-up!] This Experiment Vitamin C—whose chemical name is “ascorbic acid” and chemical formula is C6H8O6—is a water-soluble essential vitamin for humans and certain other animals, due to its involvement in a number of biological processes including wound healing and collagen synthesis (which is how it helps prevent and treat the disease called scurvy), as well as its role as an antioxidant (see https://en.wikipedia.org/wiki/Vitamin_C). Consistent with its “antioxidant” function, Vitamin C is a fairly good reducing agent, meaning that it has a reasonably good tendency to give electrons away (to oxidizing agents) and become oxidized (to form DHA, see below). Iodine, on the other hand is a reasonably good oxidizing agent, meaning it has a tendency to take electrons (from reducing agents) and become reduced (to form iodide ion, I-). In this experiment, we take advantage of the tendency of Vitamin C to undergo oxidation by aqueous iodine to determine the amount of Vitamin C in a beverage via titration. The balanced equation for the titration reaction is: C6H8O6(aq) + I2(aq) → C6H6O6(aq) + 2 H+(aq) + 2 I-(aq) (1) where C6H6O6 is the formula of a substance called “dehydroascorbic acid” (DHA, for short), the product of the oxidation of Vitamin C. Since the oxidation reaction is reversible, Vitamin C is sometimes converted into DHA in the body for transporting purposes and then re-reduced to reform Vitamin C where it is needed (e.g., inside mitochondria). For more information about DHA, which has a number of other interesting applications, you may access https://en.wikipedia.org/wiki/Dehydroascorbic_acid. Unlike in our prior titration experiment, students will not be given an assumed value for the molarity of the titrant. Rather, students will determine the molarity of the aqueous iodine titrant—a process called standardization—by reacting the iodine titrant solution with a known mass* of pure ascorbic acid. * Students will not actually be measuring out solid for each titration, as the mass of each sample would be too small to be practical. Rather, a known mass of ascorbic acid will be dissolved in water to make a solution with known volume, and then a small volume of this “stock solution” will be measured out for each titration. Knowing the initial grams of ascorbic acid and the (total) volume of stock solution prepared, in mL allows the calculation of the number of mg of ascorbic acid per mL of stock solution, so that the mg of ascorbic acid in the measured volume of each sample of solution used for each titration can be calculated. In so doing, the titration calculation process is effectively turned backwards—instead of knowing the molarity of the titrant and analyzing for an unknown amount of analyte, we know the amount of the analyte titrated and use that to determine the molarity of the titrant. More specifically, by knowing the mass of ascorbic acid used in each standardization titration, the moles of it can be calculated using the molar mass. Then the moles of iodine from the titrant (that react with it to reach the endpoint) can be calculated using the stoichiometry of the titration reaction (from the balanced equation). Since the volume of the titrant used will be determined during the titration (from the buret readings), the two quantities needed to calculate molarity—moles of iodine and volume of aqueous iodine solution (the titrant)—will both be known in the end. Once the volume is converted into liters, dividing the moles of iodine used by the volume of the iodine solution used will yield the desired molarity. The results of three determinations will be averaged, and then, with the molarity of the titrant in hand (or at least the data needed to calculate it later!), beverages that contain Vitamin C will be analyzed for their Vitamin C levels. Students may bring a beverage to analyze, but it must be clear or light in color, and have a food label that indicates the amount of Vitamin C per serving. Then, student results can be compared with information from the label to assess the accuracy of the determination/method. In any titration experiment, one must have a way to know when all of the analyte has just reacted with the amount of titrant species added. In other words, the equivalence point of a titration is supposed to be the point at which the amount of added titrant is enough to react with all of the analyte species, but not so much as to have any excess of the titrant species left over after reaction. This is what allows one to use stoichiometry to accurately calculate the amount of
Analysis of Vitamin C 2 analyte present. If you overshoot a titration trial by adding extra formula units of the titrant species, your calculations using the balanced equation will be inaccurate—the more you overshoot, the greater the inaccuracy. In practice, however, we do have to add just a tiny bit of excess titrant in order to know that all of the analyte species is gone (reacted). But how do we actually know that? We typically add some kind of indicator species to the analyte solution, which reacts with titrant once all of the analyte units are used up. The indicator typically changes color upon reaction with the titrant species, so the appearance of the color of the indicator product “indicates” to us that all of the analyte units have been reacted. We call this the endpoint of the titration. As long as the amount of excess titrant formula units is very small relative to the total amount of added titrant, the inaccuracy is typically within the level of uncertainty of the determination, and so it is considered acceptable (the difference between the endpoint and the true equivalence point is insignificant, practically speaking). In this experiment, the aqueous iodine in the titrant is, itself, colored (yellow when dilute; brown when concentrated). As such, iodine actually could function as a “built in” indicator here—when all of the Vitamin C analyte molecules have been converted into DHA, there is nothing left for the iodine to react with, and so it remains as I2 in the solution, giving the solution a faint yellow color. However, this faint yellow color is difficult to see even in an otherwise colorless solution, and is impossible to see in a sample that is already colored (e.g., common juices such as orange juice, apple juice, etc.). As such, we will take advantage of an interesting property of iodine—its tendency to bind to starch molecules and change its color to a dark blue color—to enhance the observation of the endpoint. In effect, the starch is considered an indicator in this titration, since it reacts with any excess iodine to form a product that is much more deeply colored than the unbound iodine: I2(aq) + starch → starch-I2 complex (2) (faint yellow) (dark blue) With this starch indicator present, one should be able to observe a much more precise endpoint than, for example, was obtainable in the milk of magnesia titration using the red cabbage indicator. Here, in the standardization titrations, the colorless Vitamin C solution should turn from completely colorless to a fairly dark blue color with one drop or “cut” (see videos). In addition, certain beverages (those that are not dark red or blue in color) such as orange juice, apple juice, white grape juice, Tang ®, vitamin water(s), etc., can be analyzed quite successfully for Vitamin C content using this titration reaction with starch indicator. The endpoint should yield a dark blue, almost black coloration (see videos) with again, just one drop or “cut”. However, there are limitations to the methodology described here. Dark colored beverages such as purple grape juice, cranberry juice, fruit punch, etc. cannot be analyzed this way because their existing dark color masks the observation of the blue color of the indicator product, and so the endpoint cannot be determined. Other analytical techniques (chromatography, spectrophotometry, fluorimetry, electrochemistry, etc.) would be needed to analyze such beverages.
Analysis of Vitamin C 3 Reagents and Equipment Needed • Balance • Buret, buret clamp, ring stand • Erlenmeyer flasks: 125-mL, 250-mL • Watch glasses (one large and one small) • Graduated cylinders: 10-mL, 50-mL • Vitamin C powder [ascorbic acid] (~0.2 g) • Beakers: 50-mL (1), 100-mL (2), 250-mL (1), [400-mL or 600-mL, only in rare cases] • DI water (distilled and deionized water) • Disposable (Beral) pipets (multiple) • Spatula • 50-mL volumetric flask • Dilute, aqueous iodine solution (between 0.004 and 0.006 M; precise molarity will be determined by students), ~300 mL per student (or pair) • Clear or light-colored beverage that contains some Vitamin C (e.g., apple juice, white grape juice, orange juice, Tang, other juice/drink fortified with Vitamin C, etc.) Safety: Follow all of the safety rules you agreed to when you signed your safety form. Wear safety goggles; we always wear safety goggles when doing chemistry experiments since we just don’t know what to expect. Don’t ever assume something is safe. NOTE: Aqueous iodine solutions can stain skin and clothes, so handle carefully. It is also light sensitive, so keep it away from light when not in use. Chemical Disposal: All solutions can go down the drain with flushing water.
Analysis of Vitamin C 4 Procedure, Standardization of a Solution of Aqueous Iodine Using Pure Vitamin C (Ascorbic Acid) Part 1: Preparation of Vitamin C (Ascorbic Acid) Stock Solution 1. Tare a clean and dry 100-mL beaker on the balance. Add between 0.18 and 0.20 g of ascorbic acid (Vitamin C) powder using a spatula. Record the value as precisely as the balance allows (3 or 4 significant figures, minimum). 2. Using your 50-mL graduated cylinder carefully pour ~12-15 mL of DI water into the beaker from Step 1 and swirl carefully and steadily to dissolve all of the white crystals/powder. This may take a minute or two, but all should dissolve eventually. Do not use a glass stirring rod here. Just swirl. Make sure not to spill or splash out any of the contents of the beaker. Be sure to read through the following two steps before you actually begin trying to execute them! This procedure is called a “quantitative transfer” and the goal is to get every molecule of Vitamin C (dissolved in the solution in the beaker) into the volumetric flask. So ideally, no liquid drops should be spilled or “dripped” during this process! 3. Carefully pour the Vitamin C solution from the 100-mL beaker into a clean (not necessarily dry) 50-mL volumetric flask (provided in the lab—not from your drawer), keeping the beaker inverted over the top of the volumetric flask after the pour (with the spout of the beaker physically touching the flask opening) so that no drops of solution are lost. Then, while the beaker is still inverted, carefully squirt a stream of DI water from your wash bottle onto/into the inside angled surface of the beaker just above the spout so that any drops of Vitamin C solution still in the beaker near the spout get rinsed into the 50-mL volumetric flask. 4. Remove the beaker, hold it “right-side up” in your (non-dominant) hand, and rinse the sides with several milliliters of water. Swirl the contents of the beaker a few times. Then, as in Step 3, pour this rinse into the volumetric flask, again keeping the beaker over the flask while you rinse it to assure that no drops are lost. Repeat this step (Step 4) two more times. By this process, you should have transferred all of the Vitamin C that you originally weighed out into the beaker, into the volumetric flask, and the flask should be less than half full. 5. Fill the 100-mL beaker that you just used (in Step 4) about halfway with DI water (label it as well). Carefully pour DI water into the 50-mL volumetric flask until it is about half full, and then swirl the contents of the flask well. Keep adding water, a bit at a time, and swirling after each addition, until the water level is up to the bottom of the neck of the flask only. Be very careful when you get close, because if you are still pouring when the water is filling up the neck, it will fill up really fast!! After this point, use a disposable pipet to add DI water, going dropwise once you get close to the mark/line on the neck. You want to add water only until the bottom of the meniscus is touching that mark (see images below, although these flasks are not 50-mL ones—they are larger than what you’ll be using). Note: if you do watch either of these videos, be aware that some features of their technique are not recommend by this author. From: https://www.youtube.com/watch?v=wiGUAh-_sUY From: https://youtu.be/ttkFGljBevk 6. Once the liquid level is right at the mark, cap the flask with the provided ground glass stopper, apply light pressure to the top of the stopper to make sure it does not fall out, and then carefully invert the flask (and “revert” it) 15 times (yes, you should count!) to ensure thorough mixing. Do not worry if the liquid level does not come back right to the mark at the end of this process. As long as you added water just to the mark earlier, you should record your “prepared volume of stock solution of Vitamin C” to be 50.00 mL (although the uncertainty is about 0.05 mL, not 0.01 mL). 7. Transfer the entire Vitamin C stock solution into a clean and dry 125-mL Erlenmeyer flask, and cover the top with a small watch glass. Label the flask “Vit C stock” (or put it on a paper towel labelled as such). This is your stock solution of Vitamin C; it has only been transferred out of the volumetric flask to make it easier to access with a disposable pipet in future steps.
Analysis of Vitamin C 5 Part 2: Preparation of the Sample(s) of Vitamin C (Ascorbic Acid) For Titration 8. If it’s been a bit since you’ve prepared it, swirl the 125-mL Erlenmeyer flask containing the stock solution of Vitamin C (from Step 7) to ensure it is mixed well. 9. Using a clean, dry disposable pipet (or the pipet you’ve used earlier with this solution, if this is your 2nd or 3rd titration), carefully transfer just enough of the Vitamin C solution from the 125-mL Erlenmeyer flask to a dry* 10-mL graduated cylinder (*or, if this is your 2nd or 3rd titration, just use the same graduated cylinder you used earlier, but only after rinsing it twice with ~1 mL of Vitamin C solution, each time discarding the rinse) to raise the level up to the 8.0 mL line. Read the volume as precisely as possible (to the hundredths place, if the cylinder is marked to each tenth). Record this volume in your lab notebook. This is the sample of Vitamin C for the current standardization titration. Cover the 125-mL Erlenmeyer flask with the watch glass and put it aside, together with the pipet. 10. Now do a quantitative transfer of this solution to a 250-mL Erlenmeyer flask (for titration). In other words: 10a. Carefully pour the Vitamin C solution from the 10-mL graduated cylinder into a clean (not necessarily dry) 250-mL Erlenmeyer flask, keeping the cylinder inverted over the beaker after the pour so that no drops of solution are lost. Then, while the cylinder is still inverted, squirt a short stream of DI water from your wash bottle onto/into the edge of the graduated cylinder so that any drops of Vitamin C solution still in the cylinder near the edge get rinsed into the 250-mL Erlenmeyer flask. 10b. Rinse the 10-mL graduated cylinder 3x with 1-2 milliliters of DI water from your wash bottle and add all rinses to the 250-mL Erlenmeyer flask, squirting a short stream from your wash bottle while the cylinder is over the flask, as before. After the final rinse, shake out any excess water from the graduated cylinder and put it aside for use in your 2nd and 3rd titrations. 11. Using the (water-labelled) beaker from Step 5, pour DI water into the 250-mL Erlenmeyer (titration) flask until the volume of the sample is roughly 50 mL. Do not record this volume (this is just a more convenient volume of sample for titration). 12. If this is your first titration, pour out about 10 mL of the provided starch solution [the indicator] into a (clean or DI-water-rinsed) 50-mL beaker. (If this is your 2nd or 3rd titration, you should still have leftover starch solution in the beaker from your 1st trial.) Using a clean or DI-water-rinsed disposable pipet, transfer roughly 1 mL* of this starch solution to the Vitamin C sample in the 250-mL Erlenmeyer flask. Swirl to mix. Put the flask aside (until Step 22). Keep the disposable pipet with the starch solution, and put these aside for use in preparing all subsequent titration samples. *NOTE: The line just below the bottom of the bulb on the disposable pipet marks 1.0 mL. Part 3: Preparation of the Buret 13. Attach a buret clamp to a ring stand and put aside. 14. Swirl the stock reagent bottle containing aqueous iodine (I2(aq)) [titrant solution] a few times, and then pour ~60-70 mL of it into a clean, dry, 100-mL beaker. This value need not be recorded. We are just using this smaller beaker for easier pouring of your titrant solution into the buret. Keep this solution covered when not in use. You may add more iodine solution from the stock reagent bottle into this 100-mL beaker at any time if you need more during the experiment. 15. Make three successive 2-4 mL rinses of the buret with the iodine solution (as demonstrated/done in last week’s titration experiment). In other words, for each rinse: a) Make sure the stopcock on the buret is closed (perpendicular to tube). b) Hold the buret in your non-dominant hand, at an angle, and carefully pour 2-4 mL of iodine solution into it from the 100-mL beaker (held in dominant hand). c) Turn the buret carefully on its side (so that it is almost parallel to the benchtop), rotating the body and rocking it very gently and slightly back and forth to coat the inside of the buret with the iodine solution, holding the top of the buret over a waste beaker to catch any possible spillage. d) Turn the buret back upright and drain the rinse into the waste beaker. If necessary, carefully but abruptly thrust the body downward a few inches in one swift jolt (or ask instructor to do this!) to get the flow going (the jolt removes the air bubble in the stopcock that causes liquid to remain in the buret.)
Analysis of Vitamin C 6 16. Close the stopcock, hold the buret in your non-dominant hand (at a slight angle, and carefully add iodine solution to it from the 100-mL beaker until the level is above the 0-mL mark by several mL. Place the buret in the buret clamp on the ring stand. 17. Place a waste beaker on the table/benchtop under the buret tip and adjust the height of the buret so that it is just above the top of the beaker. Do not place the waste beaker on the base of the ring stand. If necessary, rotate the ring stand/buret clamp so that the base is facing away from you, but the buret (and waste beaker) are facing toward you. The idea is to have any vessel that is under the buret sitting right on the benchtop and not on the base of the ring stand (so that it can’t tip or fall off of the base). 18. Watching the tip and stopcock area closely, open the stopcock fully to let some liquid out, pushing air out of the tip with it. If full opening doesn’t dislodge all of the air/air pockets, open it partially/slowly, and then close it off and then open again. The air may come out right away and easily, but it may also take a few iterations of this procedure. 19. Check the level of the iodine solution in the buret. If it still above the 0-mL mark, carefully open the stopcock to let some out of the buret (into the waste beaker) until the level is between 0 and 1 mL. If the level is below the 1-mL mark, carefully add iodine solution until the level is between 0 and 1 mL. 20. Cover the 100-mL beaker containing the iodine titrant solution, and put it aside. Part 4: Standardization Titration of (Three Samples of) Vitamin C (Ascorbic Acid) With the Iodine Solution 21. Make sure the buret is in proper shape to begin a titration (no leaking from tip, level is between 0 and 1 mL [at least for the 1st titration; if 2nd or 3rd titration, level may not necessarily have to be this high, but must be high enough given the expected titration volume such that the buret level will not fall below 50 mL during the titration], etc.). Record the initial buret reading in your lab notebook. 22. Place the 250-mL Erlenmeyer flask (from Step 12) directly underneath the buret tip. Adjust the height such that the buret tip is just below the top edge of the flask. 23. If this is your first standardization titration do (a) below. Otherwise, go directly to step (b) below. (a) (1st titration) Open the stopcock all the way and watch the Vitamin C solution carefully. Initially, you should see a very small faint blue region just where the titrant (I2(aq)) hits the solution which disappears almost immediately even as you continuously add titrant (see video #1 if your instructor has provided videos, though keep in mind that the flow rate may vary from that of your buret). If you do not see this blue, make sure to add the starch indicator! If you still don’t see blue, contact your instructor. As more and more titrant is added (keep stopcock opened all the way during this time!), the blue color will spread further and further into the solution, becoming a bigger and bigger patch which takes longer and longer to go away (at some point, you will close the stopcock and swirl to get the blue to go away). See videos #2 and #3 if your instructor indicates/provides. When you feel you are getting close to the endpoint, you may start adding “cuts” and assessing how quickly the color goes away to help you determine how close you are to the endpoint. (See videos #4 and #5.) When the solution remains blue after swirling (hopefully reasonably faint!), you have reached the endpoint. NOTE: The endpoint should be reached before the level of the titrant in your buret reaches 50 mL! If you are at 45 mL or lower for your buret reading and you haven’t yet reached the endpoint, contact your instructor for assistance. You ideally want to get to this point (colorless to blue) by adding one drop or one cut! But since this is the first trial, which you want to do fairly quickly to get a rough idea of the endpoint volume, you may not quite achieve this—you might overshoot it a bit. That’s okay, as long as you know that you haven’t overshot it by a ton, although if you did overshoot (even by a drop or two), you will not include the results from this trial on your report form or in your calculations. Record the final buret reading in your lab notebook and indicate whether you have overshot the titration (added more than just a cut to get the blue color to persist) or not. Also, take a picture of the sample flask (& maybe write #1 on a piece of paper towel and put the flask on the towel to label it before you take the pic). At this point, add one additional cut, just to see how much darker the solution becomes. Then do it again. You need not record your observations here, but hopefully this will give you a sense that the more you overshoot, the darker the solution gets (until it is essentially black). Go to Step 24 (skip (b) below).
Analysis of Vitamin C 7 (b) (only if doing 2nd or 3rd titration) Since you are using essentially the same volume of stock Vitamin C solution, and you are using the same iodine solution as the titrant, you should expect your total titrant volume used to be the same in trials 2 and 3 as was in trial 1, unless you overshot trial 1, in which case, it should be less than that. So, keeping that in mind, determine an “expected” total titration volume for this titration. Then, add that value to the initial buret reading for this trial. This is your “estimated final buret reading value” for this trial. Write this down somewhere (though it need not be in your notebook—if you do write it in your notebook, make sure to lab it properly, as this is not “data”!!) Now, open the stopcock fully and add iodine solution from the buret until you’re 1-2 mL “before” your estimated final buret reading. In other words, if you estimate your final reading to be 36.31 mL, then stop at about 34.5 – 35.0 mL. Swirl the flask to mix well. If you’ve calculated correctly, and your first titration was not significantly overshot, the solution should still be colorless at this point. (If not, then use this trial as your “Trial 1” and redo a new “Trial 2”.) Now proceed carefully, making only small additions and then cuts—relying on your experience from Titration #1, as well as a viewing of videos #5, #7, and #8—such that you see the transition from colorless to blue with just one cut or drop. This represents reaching the endpoint without overshooting. Record the final buret reading in your notebook, take a picture of the sample in the Erlenmeyer flask, and indicate whether you have overshot the titration or not (hopefully not, but it does happen sometimes!) 24. Discard the Erlenmeyer flask contents to waste, rinse well with tap water, discarding after each rinse, and then rinse a few times with a small amount of DI water and discard the rinses. 25. Repeat Steps 8-12 and 21-24 until you have performed three “good” (i.e., not overshot) standardization titrations of your Vitamin C stock solution, unless your instructor indicates otherwise. Procedure, Titration of a Vitamin-C-Containing Beverage of your Choice 26. If you did not bring your own beverage to analyze, choose one from the lab. Record in your lab notebook: • the specific type of beverage [not just “orange juice” or “grape juice” but brand name and type of juice as well such as “Welch’s organic white grape juice”], • the serving size, and • the amount of Vitamin C per serving. (This info can be found on the label of the beverage. You may wish to take a picture of the label as well.) 27. Swirl (but don’t shake up) the sample container of your chosen beverage. For most beverages, measure out 49-50 mL* of it in a 50-mL graduated cylinder (prerinse, if necessary). If you happen to have chosen a “100% juice” beverage, measure out 39-40 mL instead. Accurately read the volume to the nearest 0.1 mL (estimate the digit in the tenths place). Record the volume in your lab notebook. *NOTE: If this is your 2nd or 3rd titration you may wish to use 100 mL of sample if your first titration required 15 mL or less of titrant, or 30 or 40 mL of sample if your first titration required more than 50 mL of titrant. The idea here is to aim for a final “mL of titrant added” that is between 30 and 45 mL. [If your first titration required 10 mL or less, you may consider using 150 mL of sample for your 2nd titration, but you would then need to prepare your sample in a larger vessel such as a 400- or 600-mL beaker and be extra careful when swirling during the titration. If you are in this situation, ask your instructor for guidance before proceeding.] 28. Pour the sample into a clean (but not necessarily dry) 250-mL Erlenmeyer flask (or 400- or 600-mL beaker if applicable—see note above and also in Step 30 (c) below.) 29. Add roughly 1 mL of the starch solution to the sample in the 250-mL Erlenmeyer flask. Take a picture (label the sample somehow, as noted earlier). 30. Titrate as in Steps 21-23, keeping in mind the following: a) Your sample may not have all that much vitamin C present, and thus the endpoint may come unexpectedly “soon” in the 1st (rough) titration, b) The blue color will likely take longer to disappear than in the samples of pure ascorbic acid, so wait at least 30 seconds before you conclude that you are truly at the endpoint, AND
Analysis of Vitamin C 8 c) If the sample of beverage has a color and/or is not transparent, the endpoint color may look different than in the titration of the (pure) ascorbic acid samples. (It may just appear “dark” rather than blue. See videos #9, #10, and #11) NOTE: If you are concerned that your juice is a bit dark and that you’ll have trouble seeing the endpoint, you may, with your instructor’s permission, add 50 mL of DI water to your sample before starting the titration. 31. When the endpoint has been reached, make sure to record the final buret reading and indicate whether you have overshot the titration. Also take a picture. 32. Discard the flask contents to waste and rinse well with tap water and then (a bit of) DI water. 33. Repeat Steps 27-32 until you have performed three “good” (i.e., not overshot) titrations, or as your instructor directs. 34. Clean up! Before you dump out your excess vitamin C stock solution, try pouring a bit into your waste beaker and see if anything interesting happens! Can you explain what you see? 🙂 All solutions and waste contents may be flushed down the sink, ultimately.
Name: _____________________________ 9 Exp 6. Prelab: Analysis of Vitamin C 1. A sample of 0.386 g of Vitamin C (ascorbic acid) was dissolved in enough water to make a 30.0. mL (stock) solution. [NOTE: Use of chegg.com or other “homework” site to get the answers to these questions is cheating. I wrote this prelab and I do not authorize any person paid by a so-called educational website to answer these questions for students nor post the answers on the web.] How many mg of Vitamin C are in each mL of this stock solution? _______________ 2. A 6.00-mL sample of the stock solution in #1 above was used in a titration to standardize a solution of I2(aq). [NOTE: Use of chegg.com or other “homework” site to get the answers to these questions is cheating. I wrote this prelab and I do not authorize any person paid by a so-called educational website to answer these questions for students nor post the answers on the web.] The Vitamin C sample required 38.76 mL of the I2(aq) to reach the endpoint. (a) How many mg of Vitamin C were in the sample that was titrated? _______________ Hint: Use your answer to #1 above. (b) How many moles of Vitamin C were in the sample that was titrated? (MM of Vitamin C = 176.1 g/mol) _______________ (c) How many moles of I2 must have reacted with the Vitamin C in the sample? The balanced equation for the chemical reaction between Vitamin C (C6H8O6) and I2 is: C6H8O6(aq) + I2(aq) → C6H6O6(aq) + 2 H+(aq) + 2 I-(aq) _______________ (d) What must be the molarity of the solution of I2(aq) that was used as the titrant? _______________ Hint: Molarity is “moles of solute per liter of solution”. Consider how many mL of this solution were used, and how many moles of I2 were in it! 3. A 90.-mL sample of juice was titrated with the I2(aq) solution described above using a buret. The initial reading of the buret was 0.24 mL. When the endpoint was reached, the reading on the buret was 33.08 mL. How many mg of Vitamin C were in the juice sample? _______________
Name: _____________________________ 10 Exp 6. Report Form: Analysis of Vitamin C Raw Data Vitamin C Stock Solution Preparation: Mass of Vitamin C (ascorbic acid) used to prepare stock solution _______________ (Total) Volume of Vitamin C stock solution (as prepared; not how much was transferred after!) _______________ Standardization Titrations: Volume of Vitamin C (ascorbic acid) stock solution used for Titration 1 _______________ Initial iodine(aq) buret Reading, Titration 1 _______________ Final iodine(aq) buret Reading, Titration 1 _______________ Was this titration clearly (or likely) overshot? Y or N Volume of Vitamin C (ascorbic acid) stock solution used for Titration 2 _______________ Initial iodine(aq) buret Reading, Titration 2 _______________ Final iodine(aq) buret Reading, Titration 2 _______________ Was this titration clearly (or likely) overshot? Y or N Volume of Vitamin C (ascorbic acid) stock solution used for Titration 3 _______________ Initial iodine(aq) buret Reading, Titration 3 _______________ Final iodine(aq) buret Reading, Titration 3 _______________ Was this titration clearly (or likely) overshot? Y or N Titrations of Beverage of Choice Detailed description of beverage of choice: _________________________________________________________ Volume of sample of beverage (For Titration 1) _______________ Initial iodine(aq) buret Reading, Titration 1 _______________ Final iodine(aq) buret Reading, Titration 1 _______________ Was this titration clearly (or likely) overshot? Y or N Volume of sample of beverage (For Titration 2) _______________ Initial iodine(aq) buret Reading, Titration 2 _______________ Final iodine(aq) buret Reading, Titration 2 _______________ Was this titration clearly (or likely) overshot? Y or N Volume of sample of beverage (For Titration 3) _______________ Initial iodine(aq) buret Reading, Titration 3 _______________ Final iodine(aq) buret Reading, Titration 3 _______________ Was this titration clearly (or likely) overshot? Y or N
Report Form, Analysis of Vitamin C 11 Calculated Quantities and Results, Standardization # mg per mL of Vitamin C (ascorbic acid) in prepared stock solution: _______________ Titration 1 Titration 2 Titration 3 Mass of ascorbic acid used (mg) _______________ _______________ _______________ Moles of ascorbic acid used _______________ _______________ _______________ Moles I2 used _______________ _______________ _______________ Volume of I2(aq) (titrant) used (mL) _______________ _______________ _______________ Molarity of I2 in titrant solution _______________ _______________ _______________ Average molarity of I2 in titrant solution _______________ Calculated Quantities/Results, Beverage Samples Titration 1 Titration 2 Titration 3 Volume of I2(aq) (titrant) used (mL) _______________ _______________ _______________ Moles I2 used _______________ _______________ _______________ Moles of ascorbic acid in sample _______________ _______________ _______________ Mass of ascorbic acid in sample (mg) _______________ _______________ _______________ Mass of ascorbic acid per mL of sample (mg/mL)_______________ _______________ _______________ Average Mass of ascorbic acid per mL of sample (mg/mL) _______________
Report Form, Analysis of Vitamin C 12 Analysis 1. What beverage(s) did you analyze? Write down here the serving size (in mL) and the amount of Vitamin C per serving (in mg) from the label. [NOTE: If the amount of Vitamin C per serving is stated only as a percentage daily allowance (DA), contact your instructor for guidance (you may need to look up on the web the current value of DA for Vitamin C in the USA)] 2. Using the average value that you determined for the number of mg of Vitamin C per mL of beverage (calculated from your raw data), along with the volume of one serving size of your beverage (as noted in #1, from the label), calculate the (experimental) mg of Vitamin C in one serving size of your beverage. How well does your experimental value agree with the value found on the label? Comment on potential sources of error in the experiment. 3. Comment on the precision of your determinations of both (a) the molarity of the iodine solution (titrant) and (b) the mg of Vitamin C per mL of beverage. State clearly what you are looking at to assess this. 4. (a) To better quantify the precision comparison noted in #3 above, assume that a rough approximation of the % uncertainty in each of the determinations can be calculated as follows: First, calculate the range of values (for both the molarity and the mg/mL quantities) and divide it by 2 (i.e., take the largest value you got in one of your trials minus the smallest value and divide the result by 2). Then, take that result (for each quantity—molarity and mg/mL), divide it by the average value of that quantity (calculated on the previous page), and multiply by 100. Setups/Calcs for % Uncertainty of Molarity of I2 Setups/Calcs for % Uncertainty of mg/mL of Vitamin C in beverage (b) Based on this measure, which determination was more precise, the molarity of the iodine solution or the mg of Vitamin C per mL of beverage?