Question

1) What chemical is the titrant in this experiment? What chemical is the analyte in this

experiment?

Experiment #8: Measuring the Vitamin C Content of Emergen-C™ Objectives: • Students will learn how to use a burette. • Students will learn how to use the iodine starch indicator system to monitor oxidation/reduction reactions. Students will learn how to conduct a titration to determine the amount of analyte in an unknown solution. • Students will learn about ascorbic acid's role as a biological reducing agent. Introduction: Vitamin C, also known as ascorbic acid, is a molecule that serves as a biological reducing agent. It is found in many fruits and vegetables, including citrus fruits, tomatoes, strawberries, broccoli, kale, and Brussels sprouts, to name just a few. The oxidation of ascorbic acid liberates two electrons that can be delivered to a variety of substances, including biomolecules, metal ions, and iodine molecules, as is the case with this experiment. MO- GO HO--- на ЕС HO-200 + 26 2H HC- но он Ascorbic Acid 0 0 Dehydroascorbic Acid Ascorbic acid (C.H.O.) plays a key role in human health by aiding in the production of a type of connective tissue called collagen. More specifically, ascorbic acid facilitates the activity of several key enzymes that are required for collagen synthesis. Scurvy is a disease caused by insufficient collagen synthesis attributed to vitamin C deficiency. The disease is characterized by deterioration of the skin, joint pain, bleeding in the joints, and anemia. In advanced cases, the patient may develop swollen gums and experience tooth loss, as well as organ failure, convulsions, coma, and death. In addition to preventing scurvy, vitamin C also plays a role in the absorption of iron and in the deactivation of a class of potentially destructive molecules called free radicals. The role vitamin C plays in human health is multifaceted and indisputable, so it will come as no surprise that it is commonly added to beverages marketed as 'healthy alternatives to sugary carbonated drinks. Whether or not beverages like Emergen-C™ are actually, the fact remains that ascorbic acid is an essential nutrient. It is required for the normal functioning of the human body but it cannot be produced by the multitude of biochemical pathways that comprise human physiology. In this experiment, you will use an oxidation-reduction reaction to determine the amount of vitamin C in a packet of Emergen-C Super Orange (9.1 g of drink powder contains 1000 mg of ascorbic acid). The procedure you will follow, called a redox titration, involves slowly adding an oxidizing agent into a solution containing ascorbic acid. Once the oxidizing agent (the titrant) has consumed all of the ascorbic acid, a side reaction will occur that produces a color change in the solution. The moment that this color change occurs, the redox titration has reached its end point and the volume and

concentration of the titrant can be used along with the mole ratios from the reaction equations to determine the amount of ascorbic acid in the solution. Chemical Reactions involved with the redox titration: Equation 1 shows the reduction of an iodate ion in the presence of acid to form iodine: 10, +6 +5e 12 + 3 H2O Equation 2 shows the oxidation of two iodide ions to form iodine: 51-5/2 12.5 e Equation 3 shows the result of combining the reduction and oxidation reactions shown above in equations 1 and 2: 10, +6 +51 - 3 12 + 3 H20 Notice that the reduction of the iodate ion in equation 1 requires the addition of five electrons. As a result, equation was scaled by 5/2 so that this equation produces the five electrons needed to reduce the iodate ion. Equation 3 shows the summation of equations 1 and 2 and represents the balanced oxidation-reduction equation that we can use to measure the amount of vitamin C in a sample. The iodine formed in equation number 3 will be reduced immediately back into two lodide ions by the ascorbic acid, which is simultaneously oxidized into dehydroascorbic acid, as shown below in equation 4: HO-ở- t e =0 + 12 - 21° + 2H HO он Ascorbic Acid The reaction shown in equation 4 will continue as long as there is ascorbic acid present to reduce the iodine that forms. When all of the ascorbic acid has been oxidized, any additional lodine that forms cannot be reduced back into iodide ions. Instead, this iodine will persist and react with the starch indicator resulting in a blue starch-iodine complex When this complex forms, the end point of the titration has been reached and the volume of titrant that was added up to this point can then be used to determine the amount of ascorbic acid that was present in the original solution. Solutions used in this experiment: • Potassium iodate (KIO: 0.0100 M): The potassium iodate solution will be used as the titrant. It will be dispensed from a burette 0.50 ml at a time until the end point is reached. The volume (and molarity) of KIO, solution added to reach the end point is used to calculate the amount of ascorbic acid in the sample

• Potassium iodide (KI: 0.60 M): The potassium iodide solution will be added to the sample of juice drink and provides the iodide ions that are needed to react with the KIO, to form the iodine molecules that will react with the ascorbic acid. Starch indicator solution (1.0% w/v): Starch molecules form a complex with iodine that has a deep blue color (if the solutions are otherwise colorless). Once all of the ascorbic acid has been oxidized, iodine will accumulate in the solution and react with the starch producing a color change that reveals the end point of the titration. It can be assumed that the ascorbic acid in the sample has been completely oxidized at the instant the starch indicator changes color. Hydrochloric acid (1.0 M): The reduction of the 10, ion also requires six H ions. These hydrogen ions will be provided by the hydrochloric acid solution. • Chemical Hazards: This experiment utilizes chemicals that are hazardous. It is critical that each student is prepared to participate in the experiment by wearing closed-toe shoes, a shirt that covers the entire torso and shoulders, as well as long pants that cover the entire leg (including the ankle area). Students must also wear safety goggles and nitrile gloves during this experiment. Students who fail to follow these guidelines will be dismissed from lab and will forfeit all the points associated with the experiment. Experimental procedure: Part 1: Standardizing the concentration of the titrant OD 1) Use a graduated cylinder to transfer 20.00 mL of 0.0200 M ascorbic acid solution to a 125 mL Erlenmeyer flask. Use a minimal amount (3-4 mL) of deionized water to rinse the residue from the graduated cylinder into the Erlenmeyer flask. DO 2) Add 12 drops of 1.0% starch solution to the flask. OD 3) Transfer 5.00 mL of 1.0 M hydrochloric acid to the flask. DO 4) Add 1.00 mL of 0.60 M potassium iodide to the flask. OD 5) Fill the burette with the 0.0100 M potassium iodate solution to eye level. Carefully open the stopcock and allow a couple of milliliters to drain in order to fill the tip of the buret with titrant and to ensure that the meniscus is sitting exactly on a whole number volume. Before beginning the titration, make note of the graduation on which the meniscus sits so that you always know the exact volume added to the flask DO 6) Titrate the standard ascorbic acid solution by transferring the potassium iodate from the burette to the Erlenmeyer flask in 0.50 mL increments. Be sure to swirl the flask gently during each addition to afford good mixing. Titrate until the solution in the flask turns a light blue color. Make note of the color of the solution when the end point has been reached. OD 7) Calculate the experimental molarity of the titrant (the iodate solution) using the volume and molarity of the ascorbic acid solution, the mole ratios from equations 3 and 4, and the volume and molarity of the iodate solution used to reach the end point. At the end of the titration, dispose of the solution in the Erlenmeyer flask in the waste container located in the back fume hood.

8) Repeat steps 1-7 above to confirm the concentration of the titrant. Rather than simply adding the exact same amount of titrant to a second sample of ascorbic acid, focus on titrating drop by drop until the end point is reached. Be sure to define the end point using the exact same color change that was observed in your first trial. Part 2: Determine the mass of Vitamin C in a packet of Emergen-C" OD 1) Weigh out -1.5 g of Emergen-C Super Orange powdered drink mix and transfer it to a 125 mL Erlenmeyer flask. Be careful to transfer every speck of the Emergen-C™ powder from the weighing paper to the flask. Accurately record the mass on the data sheet and be sure to record all of the digits displayed on the balance. DO 2) Dissolve the powder by adding -20 mL of deionized water to the flask. Swirl it gently until the solid dissolves. 003) Add 12 drops of 1.0% starch solution to the flask. DO 4) Add 5.00 mL of 1.0 M HCl solution to the same flask used in steps 1-3 above. 005) Lastly, transfer 1.00 mL of 0.600 M potassium iodide to the flask. DO 6) Before you begin the titration, you may need to top off the burette with additional titrant solution. 007) Titrate the Emergen-C solution by transferring the potassium iodate from the burette directly to the Erlenmeyer flask in 0.50 mL increments. Swirl the flask gently during each addition to effect good mixing. Continue until the solution in the flask turns a brownish blue color. OD 8) Calculate the amount of the ascorbic acid in the Emergen-C™ solution using the volume of iodate solution used to reach the end point and the average experimental molarity of the iodate solution determined in part 1. 9) Repeat steps 1-8 above to confirm the amount of ascorbic acid in the Emergen-C™ packet. Rather than simply adding the exact same amount of titrant to a second sample of ascorbic acid, your second titration should be conducted by adding the iodate solution 0.50 mL at a time until you are within -1-2 mL of the predicted endpoint. Then the titrant should be added drop-by-drop until the desired color change has been achieved. Be sure to define the end point using the exact same color change that was observed in your first trial. At the end of each titration, dispose of the solution in the Erlenmeyer flask in the waste container located in the back fume hood. Pre-lab Questions: 1) What chemical is the titrant in this experiment? What chemical is the analyte in this experiment? (2 points) 2) What chemical reaction does the titrant contribute to this experiment? (1 point) 3) Use the mole ratios contained in both equation 3 and equation 4 to calculate the volume in mL) of 0.0100 M potassium iodate solution needed to completely oxidize all of the ascorbic acid contained in 20.00 mL of 0.0200 M ascorbic acid solution. (1 point)

Answer 1

The experiment is determination of amorent of ascorbic acid in unknown Vitamin-C sample. This has been done by titrating it against potassium iodate solution. This Titrant : KIO₃ (Potassium ioclate) Analyte : Ascorbic acid