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Question: FIRST ORDER REACTION: DECOMPOSITION OF HYDROGEN PEROXIDE EXPERIMENT WAS PERFORMEND. 100 ML 1 M OF PEROXIDE solution WAS MIXED WITH 10 ML of 1% Kl solution WITH ACTED AS A CATALYST.The progress of the reaction is now followed by titrating 10 cm3 samples at approximately 10 minutes intervals. When about 9 minutes have elapsed, a 10 cm3 sample of the reaction mixture WAS TRANSFERED by means of the 10 cm3 measuring cylinder into the 54 Erlenmeyer, 5 cm3 of dilute sulfuric acid ADDED, and titrated rapidly with permanganate. the exact time at which the first drop of the permanganate is added was recorded. In as much as the peroxide continues to decompose until all of it has reacted with permanganate, the titration should be continued as rapidly as possible. the volume of permanganate required was recorded. Titrated successively 10 cm3 samples of the reaction mixture at approximately 10 minutes interval until 8 samples have been analyzed. from the analysis peroxide concetration decreased form 1 M to 0.564 M . velocity constant was found to be 7.63×10^(-3) form the slope of the graph and 9.09×10^(-3) from calculations. half life was found to be at 76.26 minutes .discuss
From your experiment, it seems you've successfully displayed the decomposition of hydrogen peroxide (H2O2) in a first-order reaction. The catalyst used (KI or potassium iodide) increases the rate of the reaction without getting consumed itself. The decrease in peroxide concentration from 1M to 0.564M throughout the experiment signifies the gradual decomposition of hydrogen peroxide. The rate constant (k) of the reaction was determined from both the graph slope and calculation, which was relatively close (7.63x10^-3 and 9.09x10^-3 respectively). These values of k suggest the reaction proceeds at a moderate pace. The concept of the half-life of a reaction is very important in first-order kinetics. It is the time taken for the concentration of the reactant to decrease to half of its initial concentration. In this case, that was found to be 76.26 minutes, which suggests that hydrogen peroxide decomposes at a relatively slow rate in your experiment. In first-order kinetics, the rate of a reaction is directly proportional to the concentration of one reactant. This is evidenced in your experiment by the gradual decrease in peroxide concentration over time. It's important to also consider possible sources of errors while evaluating the results, such as the volume measurement errors, uneven catalyst distribution, effects of temperature changes, the accuracy of timing when taking samples, and the possible continued reaction of peroxide while titration is happening. Conclusively, your data seems to align well with the theory of first-order reaction kinetics, more specifically the kinetics of the decomposition of hydrogen peroxide. Always aim to minimize errors for meticulous experiment execution and accurate results.
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