Chemical Kinetics Chemistry Lab Report
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Chemistry Lab Report: Chemical Kinetics
Procedure
The first part involved determining the relationship between reactants’ concentration and the speed of reaction. A mixture of KI and Na2S2O3, H2O was placed in a reaction flask labeled “Reaction Flask 1”. Another mixture of KBrO3 and HC1 was placed in a second flask labeled “Reaction Flask 2”. After a thorough mixing of the contents of each flask, the mixtures were combined by pouring the contents of the second reaction flask into the first reaction flask. The contents were agitated by stirring until the solution turned blue. A stop watch was used to determine the time it took for a blue coloration to appear. The temperature of the mixture at the point of turning blue was recorded using a thermometer. Five mixtures with different concentrations of the reagents were used for the experiment.
The second part involved investigating the influence of temperature on the rate of reaction. 10ml of 0.010M KI was mixed with 0.0010M Na2S2O3 in reaction flask one and 10ml of distilled water was added. 10ml of 0.040M KBrO3 was mixed with 10ml 0.100M HCL and 3 ml of starch suspension in Reaction Flask 2. The solutions were cooled to 100 c and then mixed. Time taken for a blue coloration to appear was recorded and the procedures repeated at temperatures of 0 and 400 c.
The third part involved investigating the influence of a catalyst on the rate of reaction. Mixing of the reagents was done as in the second part, and a drop of 0.5 M (NH4)2MoO4 (a catalyst) was added to Reaction Flask 2. Time taken for the appearance of a blue coloration in both flasks was recorded.
Results and Discussion
The recorded temperatures for the first part of the experiment were 20.80 c, which were rounded-off to 210 c. The rate of reaction was noted to increase as the concentration of the reactants was increased. It was observed that the concentration of different ions affected the overall results to varying extents. Doubling the concentration of iodide ions increased the rate of reaction by a factor of two. Also, doubling the concentration of bromate ions halved the overall time taken for the reaction. On the other hand, doubling the concentration of hydrogen ions resulted in a three-fold increase in the rate of reaction. The pattern of variation in the speed of reaction at different concentrations was observed in all trials performed. The order of reaction for [I-] and [BrO3-] was 1 while that of [H+] was 2 as obtained from the trials. Hence, the overall value for the order of reaction was 4. The observation concurred with the kinetic theory of reaction which states that the rate of a reaction is directly proportional to the concentration of the reagents involved. The theory explains the phenomenon by indicating that increasing the concentration of reactants results in a high number of molecules. Chances of interactions between the reacting molecules increase as the number of the particles increases. Reactions that have higher number of collisions at a given time are faster than the ones with a low number.
The recorded times for the reactions at temperatures of 40, 20.8, 10, and 1.80 c were 60, 160, 406, and 660 sec respectively. The rate of reaction was also found to increase with an increase in temperature. The reaction took a shorter time at higher temperatures than it did at lower temperatures. The observation was in accordance with the kinetic theory of reactions. The theory suggests that when particles get heated, they acquire kinetic energy and they make more movements resulting in more collisions that account for increased chances of reacting. At cold temperatures, particles have low energy and they tend to remain immobile. As a result, there are reduced chances of collisions between particles. Product formation only occurs after the interaction of reagents’ particles. At low temperatures, only a few particles collide in a given time. Therefore, the rate of reaction and product generation is slow compared to a time when there are more collisions. It is important to note that all particles may eventually react even at low temperature, but the process would take a considerably longer time. The rate of reaction is the inverse of time taken, and therefore, reactions that take a long time occur at a slow rate. As observed in the experiment, a slight change in temperature may translate into a significantly large variation in the rate of reaction. It is possible to determine the energy of activation for a particular reaction by plotting the rate constant against the inverse of time (rate of reaction) taken for the reaction to occur. It is expected that the rate constant would increase with temperature. The occurrence is in accordance with Arrhenius behavior which suggests that a high value of activation energy would mean a high correlation between temperature and the rate constant. The conventional energy of activation for a clock reaction for iodine is 54 KJMol-1. The value obtained from the experiment was 45.3 KJMol-1. Closeness to the value in the obtained results depends on the level of accuracy involved in the experiment. As obtained from calculations, the equation for a graph of the natural log of the rate constant against the rate of reaction would have the equation y= -5526x + 26.71.
15 seconds were recorded for the catalyzed reaction compared to 174 seconds recorded in the absence of the catalyst. The use of a catalyst was found to have a great impact on the speed of reaction. The rate of reaction in the presence of a catalyst was eleven times faster than it was in the absence of the catalyst. Catalysts affect the speed of reaction by decreasing the activation energy required to initiate reactions. They also offer a surface on which reactions can take place. Catalysts are never used up in reactions, and they only create new transition states through which reactions would proceed. Even small amounts of catalysts would have a significant impact on the speed of reactions.
Conclusion
The three parts of the experiment investigated factors that influence the rate of chemical reactions. Among them are temperature, reagents’ concentration, and the presence of a catalyst. One can manipulate the speed of reactions by varying either of the factors. When studying the influence of a particular factor on the rate of reaction, it is important to hold other variables constant. The procedure would ensure that variations observed are specifically as a result of the factor of interest. A combination of factors such as high temperature, high concentration of the reagents, and the presence of a catalyst would result in fast reactions. The experiment revealed the expected results as findings correlated with the theory of kinetics. The obtained value for activation energy was 45.3 KJmol-1 and that of the rate constant was 2929 1/M3s. the order of reaction was 4. The results were credible and reliable. However, there could have been improvements in the experiment to enhance accuracy and precision. Among them include ensuring that the reagents used were free from contamination, and they were appropriate for use. For instance, ensuring that the starch used in the experiment was fresh would have been a recommendable practice.
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