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Chemical Kinetics

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Chemistry Lab Report: Chemical Kinetics

Chemistry Lab Report: Chemical Kinetics



The first part involved determining the relationship between reactants’ concentration and
the speed of reaction. A mixture of KI and Na 2 S 2 O 3, H 2 O was placed in a reaction flask labeled
“Reaction Flask 1”. Another mixture of KBrO 3 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 Na 2 S 2 O 3 in reaction flask one and 10ml of distilled
water was added. 10ml of 0.040M KBrO 3 was mixed with 10ml 0.100M HCL and 3 ml of starch
suspension in Reaction Flask 2. The solutions were cooled to 10 0 c and then mixed. Time taken
for a blue coloration to appear was recorded and the procedures repeated at temperatures of 0 and
40 0 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 (NH 4 ) 2 MoO 4 (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 20.8 0 c, which were
rounded-off to 21 0 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 [BrO 3 -] 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 1.8 0 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.



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/M 3 s. 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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