- Time Trial One 587 Seconds Trial Two 1085 Seconds Trial Three 419 Seconds Trial Four 1316 Seconds 3 Calculate The R 1 (1.3 MiB) Viewed 34 times
Method Rinse all glassware required for this experiment with both tap and distilled water before you begin! Any piece of equipment might still have residue left from the last time this experiment was run and it will ruin your results! Also, you MUST carefully rinse your graduated cylinders after each use to prevent contamination! If you do not, you will have to do the entire experiment over again! Be sure that each drop of liquid you measure gets quantitatively transferred to your reaction flask (by holding the graduated cylinder upside down for 15 seconds) before going on to the next measurement. Label all containers! Be organized! Check off each measurement as you make it! Fill a 600 mL beaker with distilled water, label it, and use it to make the following three solutions. Use a clean, dry, graduated cylinder to make measurements. Buffered Iodide Solution To a 400 mL beaker, add: 25.0 mL acetic acid/sodium acetate buffer 6.0 mL of 0.5% starch solution 6.0 mL of 1.0 M KI then fill to the 250 mL mark with distilled water. Label the beaker. 0.025 M Thiosulfate Solution To a 250 mL beaker, add: Peroxide Solution To a 250 mL beaker, add: 30.0 mL of 1.0 M H₂O₂ then fill to the 150 mL mark with distilled water. Label the beaker. Mix each solution with a clean, dry, glass stirring rod. Use these solutions for the rest of the experiment! Do not go back to the stock solutions! Label four clean, dry, 125 mL Erlenmeyer flasks with the numbers 1 through 4. Following Table 1, transfer the required volume of each solution into each flask. Consider measuring one solution for all four flasks before rinsing and draining the cylinder. Record volumes to the nearest 0.1 mL. Table 1 Buffered Iodide Solution (mL) Thiosulfate Solution (mL) Distilled Water (mL) #1 #2 #3 #4 40.0 40.0 40.0 40.0 10.0 10.0 10.0 5.0 10.0 20.0 0.0 5.0 Put flask #4 in a 1000 mL beaker containing ice and water. Measure 30 mL of peroxide solution into a clean, dry, 50 mL Erlenmeyer flask and add this container to the ice as well. Continue with the room temperature rate measurements while these solutions cool. Measure the temperature of each flask before beginning the reaction. The amount of peroxide solution to transfer to each of the flasks is given in Table 2. Begin timing as soon as you add the peroxide to the flask. Swirl the flask contents periodically. Record the start time as well as the time at the 5.0 mL of 1.0 M Na2S₂O3 then fill to the 200 mL mark with distilled first appearance of a blue colour in the flask to water. Label the beaker. the nearest second.
Table 2 Peroxide Solution (mL) #1 20.0 10.0 30.0 #2 #3 #4 30.0 If you use a clock to time your reactions, record the hour, minute, and second for both times. Do not try to count the sweeps of the second hand! When the two flasks on ice have cooled to below 3°C (rinse and dry the thermometer between measurements!) add the peroxide in the small flask to flask #4 and time the rate of reaction as before. Keep the reaction mixture on ice! You may remove the flask briefly to swirl the contents periodically. Measure the final temperature of each flask before pouring solutions down the drain with plenty of water. Details of the Reaction When peroxide is added to the reaction flask, it immediately begins to oxidize the iodide anions and produces iodine molecules (E1). This reaction continues until one of the reactants is used up. At the same time, the thiosulfate in the flask reacts with these iodine molecules before they have a chance to combine with the starch and converts them back into iodide anions (E2). Once the thiosulfate is used up, the iodine molecules produced react with the starch and the solution turns blue. The amount of thiosulfate added to the flask, therefore, determines the length of time the reaction can continue before the blue colour is observed. NOTICE! The same amount of peroxide reacted in each of the room temperature trials (since each flask contained the same amount of thiosulfate) but the time it took depended on the initial amount of peroxide. The higher the initial concentration of peroxide, the faster the reaction. Since [H] and remain constant during the timing period, and the amount of H₂O₂ reacted is negligible compared to the initial amounts of H₂O2, it would appear that the variable causing the different reaction speeds is the differing initial concentrations of peroxide. You will be asked to calculate the rate of reaction, the order of reaction, the rate constant, and the activation energy for the oxidation of iodide by hydrogen peroxide. Calculating the Rate of Reaction Based on the details of this reaction as discussed above, the rate of reaction is the change in peroxide concentration per second: E6: rate= Δ[Η,0,] time You can calculate the change in peroxide concentration from the amount of sodium thiosulfate added. For every mole of peroxide reacted in E1, one mole of iodine is produced, which reacts with two moles of thiosulfate. Use dimensional analysis, the molarity, and the volume of the beaker sodium thiosulfate solution used to calculate moles of thiosulfate reacted. Divide by 2 to get the moles of peroxide reacted in the measured time period. Divide by the total reaction volume to get A[H₂O₂], then divide by total reaction time in seconds to get the rate in units of mol.L¹.s¹. Calculating the Order of Reaction If the order of a reaction with respect to a given reactant is zero, the rate does not depend on the concentration of that reactant and will stay the same no matter what the concentration of the reactant is. If the order is one, doubling the concentration doubles the rate; if the order is two, doubling the concentration quadruples