Iodine Clock Reaction - Critical Skills for Biomedical Science PDF

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The iodine clock reaction is a classical chemical clock demonstration experiment to display chemical kinetics in action. I had to explain the effects of concentration on the rate of the reaction in the experiment I carried out in the lab....

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Observing the reformation of the blue/black colour with a specific target time Abstract: Over the years, clock reactions have been a method of studying chemical kinetics and are an example of an initial rate process (no name, 2019). The rate can be calculated according to the time taken for an observed endpoint to happen and this is a colour shift of the blue and black iodine clock reaction. Changing the concentration of a reactant whilst ensuring others are kept constant can result in different times for colour changes to take place. This experiment indicated that increasing the concentration of ascorbic acid will result in an increase in the time of the reaction for the colourless solution to turn back to blue-black, however, the results obtained doesn’t match the expectation of the hypothesis. Introduction: The iodine clock reaction is particularly well-suited as an introduction to reaction rates and kinetics. In clock reactions, they contain a complex mixture of chemicals which react in order to cause a physical change after a period of time (Faculty.sites, 2015). The term clock reaction comes from the time taken for a quantity of product to form changes when the concentration of one of the reactants varies, due to different physical changes occurring at different time intervals. Usually, there is a noticeable endpoint, such as a colour change that indicates when the desired product has been formed. In this case from colourless back to blue-black colour. In the iodine clock reaction, the triiodide ion reacts with the starch in order to form a blue-black colour which is shown in the following equation: I₃⁻ + starch→blue-black complex To generate triiodide ions, the potassium iodide solution must react with hydrogen peroxide solution, the triiodide ion cannot form the iodine blue-black complex whilst this reaction is occurring and therefore starch is added to the solution. This is shown in the equation below: 3 I⁻ + 2 H⁺ + H₂O₂ →I₃⁻ + 2 H₂O (The iodine was from a potassium iodide

solution). The hypothesis of this experiment was as the concentration of ascorbic acid increases, the time of the reaction also increases. The aim of this practical was to determine an experimental design for the iodine clock reaction to occur using a group-based method such as altering ascorbic acid volumes to convert the triiodide ion back into the non-complex forming iodide ion in a specified amount of time (47 seconds), this was to be specified by a colour change from colourless to blue-black. Method: The following reagents were given to work with: 10 mL 3% hydrogen peroxide

solution, 10 ml 1% potassium iodide solution, 5 mL 50 mM ascorbic acid, 5 mL starch solution, and deionised water. 20 test tubes were also provided, and a stopwatch. A P1000 micropipette and several blue tips were also required to dispense the different solutions, and a P200 micropipette with several yellow tips to dispense the ascorbic acid. Graph paper was given to construct a suitable graph with respect to the relationship between the two variables once the results were obtained. It was important to ensure that all of the test tubes added up to 1000 μL; this includes 250 μLof hydrogen peroxide, 250 μL of potassium iodide and 250 μL of starch. The concentration of ascorbic acid diluted using water was as shown in figure one. Firstly 250 μLof the potassium iodide solution using a P1000 micropipette was dispensed into test tube “A”. Using figure one 0 μLof ascorbic acid and 250 μLof deionised water was dispensed into test tube A. In a separate test tube "B", 250 μLof the hydrogen peroxide solution was mixed with 250 μLof starch using a P1000 micropipette but ensured that new tips were used for each solution. Using a P1000 micropipette the 500 μLof the solution in test tube B was transferred into test tube A, the stopwatch was started as soon as the solution changed from colourless to blueblack and results were recorded in a suitable table. These steps were repeated again using different test tubes whilst ensuring the ascorbic acid increased by 10 each time (30, 40, 50, 60, 70) and the volume of water changed accordingly as shown by figure one so that both add up to a final volume of 250 ul. Three trials were carried out for each concentration. Once the 47 seconds was achieved in trial 2 a lab technician observed trial 3. A graph was then generated using the mean calculated to look at the independent and dependent variable and therefore do a regression to determine at what concentration or what volume of ascorbic acid would give the colour change. Figure 1: Dilutions of 50mM ascorbic acid

Figure one displays the dilution of ascorbic acid calculated using c1v1=c2v2. The concentration of ascorbic acid had to be diluted because the volume of water was changed so that it was kept constant to 250 ml. Results: Figure 2: Time taken for solution to turn blue-black (seconds) over concentration of ascorbic acid (mM)

The independent variable (x-axis) was the concentration of ascorbic acid, and the dependent variable (y-axis) was the time taken for the solution to turn back to a blue-black colour. As the concentration of ascorbic acid increased, the time taken for the solution to turn blue-black also increased. For example, at a concentration of 6 (mM) the average time taken for solution to turn blue-black was 24.87 (s), this increased to 47.70 (s) at a concentration of 8 (mM). The results were very inconsistent between 10 -14 concentration of ascorbic acid. They display the 4 anomalies which were found in the results obtained. The average time taken for solution to turn blue-black increased between 6-8 concentration of ascorbic acid but then at ascorbic acid concentration of 8 mM, this decreased from 47.70 (s) to 35.97(s) at a concentration of 10 (mM). Then at a concentration of 12 (mM) the average time taken went back up and increased drastically to 60.00(s), this did not fit the pattern. Using c1v1=c2v2 the concentration of ascorbic acid needed at 47 seconds was calculated. c1was the initial concentration of ascorbic acid (50mM). c2=11.44; the calculation y=4.1075x was used to calculate the amount of ascorbic acid needed at 47 seconds, so 47 was divided by 4.1075 which gave ascorbic concentration of 11.44 mM. v2= the overall final volume of both ascorbic acid and water which will always remain as 250 ml. Therefore, by rearranging the equation: v1 = 11.44 x 250 = 2860 2860/50 = 57.21 μL This means 57.21 μLof ascorbic acid is required for the reaction to turn from

colourless to blue-black in 47 seconds. 192.79 μLof water was needed because the final volume is 250 μL. The r2 value as shown in the graph is 0.7437, this means there is a strong positive correlation as it is closer to 1. Figure 3:

This shows the results that were obtained during all 3 trials for the different sets of concentration of ascorbic acid. The concentration of ascorbic acid at 6mM had the quickest time, an average time of 24.87(s), for the solution to change from colourless to blue-black colour. On the other hand, the longest time it took for colour change was 60.00(s) with an ascorbic acid concentration of 12mM. Discussion and Conclusion: Based on the results obtained from this experiment, the hypothesis has not been supported, the increase that took place was only between ascorbic acid concentrations of 6-8 mM as shown in figure 3, after that the results obtained between ascorbic acid concentrations of 10-14 are very inconsistent and they don’t fit the pattern which are anomalies as shown in the graph. The graph also calculated that for 47 seconds to be the time for blue-black colour to return, the concentration of ascorbic acid had to be approximately 57.21μL. 192.79μL of water was needed because the final volume is 250 μL. As long as there is any ascorbic acid present, the triiodide ions will be reduced back to iodide ions (Wright, S. 2002). Once all of the ascorbic acid is used up, the triiodide ion accumulates and displays a colourless change to the blue-black (Shakashiei,B.1992). This is why as shown in figure 3, the results between 6-8 mM ascorbic acid concentration, less time is being taken for colour change to occur when the concentration of ascorbic acid is lower, due to there being less ascorbic acid to reduce the triiodide ions to iodine to form a blue-black complex. The moment there is no triiodide (the equilibrium has shifted to the right) the colour disappears. The iodide ion and dehydroascorbic acid are the predominant items in this equation and therefore the equilibrium will shift back to the left because there is a greater concentration of the iodide and dehydroascorbic acid. This is displayed in the following equation: Ascorbic acid + I3 - Dehydroascorbic acid + 3 I+ 2 H+ (Ascorbic acid becomes oxidised to dehydroascorbic acid)

There are several limitations and reasons why the anomalies may have occurred; firstly, the time taken for the clear solution to turn back to a blue-black colour could be impacted by human error as this was measured using a stopwatch, such as starting the stopwatch too late or too early which can result in producing inaccurate results. Another limitation which could have influenced the anomalies and inconsistency of the results could be due to human error through incorrect pipetting techniques which can affect the overall volume of the solution in the test tube. Another issue could be air bubbles can be drawn into the pipette, if this occurs then the bubbles can have an influence on the optical density values and results which may lead to inaccurate results. Furthermore, cross-contamination of solutions could have also influenced the anomalies as the same tip could have been used for more than one solution, this affects the accuracy of the results. In addition, if this experiment was repeated in the future, it would be better to conduct some extra background research on the theory behind iodine clock reactions, this would be beneficial and provide a greater context which would have allowed the experiment to run smoothly. Reflection: Overall, I have greatly benefited from this experimental design as several skills were gained. Using scientific theories and explanations along with doing background research was a great skill I established during this experiment because I was able to develop a hypothesis. Moreover, during the experiment I had to measure out specific volumes of solutions accurately using the two different micropipettes, this increased my confidence in measuring accurate volumes which would have been prominent throughout the course in other practicals. It is important to carry out experiments suitably using the correct equipment, ensuring the accuracy of measurements and also ensuring health and safety are considered at all times. This is another great skill which I have obtained through this experiment as I had to read the proformas thoroughly prior to the experiment. Another great strength of this experimental design was being able to work in a professional manner with my group members, listening to each other's ideas and respecting each other’s opinions prior to the experiment hence why the experiment was successful in terms of finishing it on time and carrying the task out smoothly without any issues. A weakness I hope to refine in the future is my pipetting technique, I was not dispensing the solutions correctly into the tubes which meant I may have left some solution in the tips, this could have been another reason why anomalies were found

in the experiment. In conclusion, the overall success behind this experimental design experiment was down to the fact that it was conducted in a group setting and that majority of the skills that I applied during the practical were applied since I had carried out background research prior to the experiment and watched videos to get a better understanding of the experiment.

References Faculty.sites.uci.edu. (2015). Clock Reactions [online] Available at: http://faculty.sites.uci.edu/chem2l/files/2011/04/A01MANClockRxn.pdf [Accessed 5 Dec. 2018]. Umanitoba.ca. (2015). HYDROGEN PEROXIDE IODINE CLOCK. [online] Available at: http://www.umanitoba.ca/outreach/crystal/resources%20for%20teachers/A %20Clock%20Reaction%20C12-3-05%20&%2009.doc [Accessed 16 Dec. 2019]. Chem.libretexts.org. (2017). 1: Chemical Kinetics - The Method of Initial Rates

(Experiment) - Chemistry LibreTexts. [online] Available at: https://chem.libretexts.org/Bookshelves/Ancillary_Materials/Laboratory_Experiments/ Wet_Lab_Experiments/General_Chemistry_Labs/Online_Chemistry_Lab_Manual/Che m_12_Experiments/01%3A_Chemical_Kinetics__The_Method_of_Initial_Rates_(Experiment) [Accessed 13 Dec. 2019]. Chemdemos.uoregon.edu. (2019). Orange Juice Clock Reaction | Chemdemos. [online] Available at: https://chemdemos.uoregon.edu/demos/Orange-JuiceClock-Reaction [Accessed 14 Dec. 2019]. RSC Education. (2015). Iodine clock reaction demonstration method. [online] Available at: https://edu.rsc.org/resources/iodine-clock-reaction/744.article [Accessed 11 Dec. 2019]. Imaginationstationtoledo.org. (2012). Iodine Clock Reaction. [online] Available at: https://www.imaginationstationtoledo.org/educator/activities/iodine-clock-reaction [Accessed 11 Dec. 2019]....