Formal Report – Kinetics of Reaction: the Iodine Clock Reaction Essay Sample
- Word count: 3517
- Category: reaction
Get Full Essay
Get access to this section to get all help you need with your essay and educational issues.Get Access
Formal Report – Kinetics of Reaction: the Iodine Clock Reaction Essay Sample
In everyday life, several reactions are encountered, but still knowledge on how fast these occur and the factors affecting it were still insufficient. This study aimed to determine the different factors affecting the rate of reaction and how these factors affected it. An experiment named iodine clock reaction was done to answer the questions raised. In this study the reaction of iodide ion and peroxodisulfate (VI) ion was analyzed with the help of thiocyanate ion. The experiment was divided into three parts. First, second, and third parts of the experiment were used to examine the effects of the reactant concentrations, temperature and the presence of catalyst to the rate of reaction, respectively. The time before the reaction proceeded was recorded and calculations were done to obtain necessary information. It was identified that increasing the reactant concentration would increase the rate of reaction. Likewise, higher temperature increased the rate. Finally, the presence of catalyst had enormous effect on the reaction rate. It increased as the catalyst was added. Based on the results obtained, it can be concluded that said factors could increase the rate of reaction by increasing the temperature or reactant concentration or by adding a catalyst on the mixture.
Until now it has been thought that in chemical reactions reactants combine with one another to form products, and that we can use a balanced equation in determining the amount of products formed. But this simple explanation does not tell us the rate of the reaction. Rate of reaction refers to the speed at which reactants break down and recombine to form new substances. In order for a reaction to proceed, the reactant particles must collide and this is what so called the collision theory of chemical reactions. In addition to this theory, collision itself will not make a reaction to happen. Molecules should be reactive and must be in their proper orientation. Moreover, enough energy must also be available in order to break the chemical bonds and the minimum amount of energy required to react is termed as the activation energy, Ea.
Why do fish get spoiled much slower inside the freezer than they do in markets? Rate of reactions differ from one another. Some reactions are so fast that they are already done as soon as the reactants are combined. Some may take minutes, days, months, and even years. What causes this wide range of reaction rates? Several factors are influencing the speed of a reaction: concentration, temperature, presence of catalyst, and the physical state or nature of reactants involved. In order to study the effect of those factors, the rate law or rate equation, which is the focal part of any kinetics study, will be of great importance. Rate law has this form: Rate = k[A]m[B]n
[A] and [B] are concentrations of the reactants expressed in molarity. Aside from that, there are also reaction orders, m and n, which tell us the effect of reactant concentration to the rate. k is the rate constant which is specific for a given reaction at a given /temperature. Always remember that data of the rate law – rate, reaction orders, and rate constant – should be obtained experimentally. In this study an experiment called the iodine clock rotation was done in order to obtain those components which led the students in determining the effects of most of the factors said. Furthermore, rate law alone was insufficient in identifying the effect of temperature on the rate of the reaction. Thus, an equation known as the Arrhenius equation was used. Arrhenius equation has this form: k= Ae-Ea/RT
where k is the rate constant, A is the Arrhenius constant, e is the base of natural logarithms, T is the temperature expressed in K, Ea is the activation energy and R is the universal gas constant (8.314 J•mol-1•°K-1). Thorough explanation will be discussed in the methodology.
The objectives of this study were to explain the kinetics of the reaction of Iodide ion and Peroxodisulfate (VI) ion, to determine the rate law of the reaction, identify the effects of concentration of the reactants, temperature, and presence of catalyst to the reaction rate. This experiment could provide even an ordinary person an explanation to several chemical reactions happening in his everyday life: why does milk turn sour more slowly in the refrigerator than it does at room temperature, why does wood burn faster in pure oxygen than in air. It may also supply basic knowledge about the processes and chemical reactions happening inside our body.
Experimental reagents prepared for this experiment were potassium iodide (KI), potassium Peroxodisulfate (K2S2O8), starch, potassium chloride (KCl), potassium thiosulfate (K2S2O3), potassium sulfate (K2SO4), and cupric sulfate (CuSO4). The glasswares, materials, and equipment used were indicated in Table 1 of Appendix A. The experiment was divided into 3 parts. The first, second, and third parts were used to determine the effects of concentration of the reactants, temperature and catalyst, respectively, on the reaction rate. On the first part, 2 beakers with different contents were prepared. 5 runs were prepared, and each run had different volumes of the substances.
The different runs for the effect of Peroxodisulfate (VI) and iodide ion concentrations on reaction rate were presented on Appendix A. Each run had different concentrations of substances to test the effect of the change in amount of concentration to the rate of the reaction. Contents of beaker A and B were mixed; the timer was started immediately, and was stopped once a blue color appeared in the mixture.  Once Peroxodisulfate (VI) ions and iodide ions combine they produce I2 molecules which in turn reacted with the starch added to form a blue color. I2 reacted with the starch as fast as it was produced. Hence, it was very difficult to measure the rate of its reaction.
So as to address the problem, S2O3 2- ion was added. This ion destroyed the I2 by reducing it back to I- as fast as it was produced. The amount of S2O32- ion added was just small so that it would not consume so much time. If the initial concentration of S2O32- ion was kept very small, Δ[S2O32-] would be small and Δ[S2O82-] would be even smaller. Consequently, there would be little change in the concentration of the reactants during the elapsed time Δt. This was a necessary condition for the initial rates method. The rate of the reaction was also affected by degree by the overall concentration of ions and their charges. Therefore, ionic substances that will not react with the other substances must be added to have a consistent concentration of the ions and the charges. In short they were added in order to maintain the ionic strength in the solutions. In this experiment those ionic substances were KCl and K2SO4.
On the second part of the experiment, the amount of concentration in run 2 was used. It was further divided into two subparts. On the first subpart, contents of the beakers were heated to 50 0C (323 K) using a hot plate. On the other hand, the same contents of the other beakers had reached 5 0C (278 K) by using an ice bath. The timer was started once contact was made and was stopped when blue color appeared. The time was recorded afterwards. This was done to study the effect of temperature on the speed of the reaction.
The last part focused on the effect of a catalyst on the reaction rate. This time CuSO4 was used as a catalyst. Contents of Run 2 were prepared again. Beakers A and B were combined. Then, 4 drops of 1 M CuSO4 was immediately poured after contents of beaker A was added to the contents of beaker B. The reaction was timed until the mixture turns blue.
RESULTS AND DISCUSSION
In order to know the effect of each factor, the rate of the reaction must be known. Formula (1) Rate= ∆[A]∆t
The concentrations of S2O82-, S2O32- , I- were calculated using the equation:
Formula (2) M1V1 = M2V2
The concentrations of the following ions in the mixture among different runs were presented in Table 1. More information was provided in Table 2 of Appendix A. Runs| [S2O82-]| [I-]| [S2O32-]| Reaction time, s| Rate, M s-1| 1| 0.02| 0.08| 0.0008| 42.7| 3.4 x 10-6|
2| 0.02| 0.04| 0.0008| 93.81| 4.3 x 10-6|
3| 0.02| 0.02| 0.0008| 142.33| 2.8 x 10-6|
4| 0.03| 0.04| 0.0008| 55.78| 7.2 x 10-6|
5| 0.04| 0.04| 0.0008| 43.49| 9.2 x 10-6|
Table 1. Effect of Reactant Concentration on Reaction Rate
Sample calculations were provided in sample calculation 1 of Appendix B. In this experiment, potassium Peroxodisulfate (K2S2O8) was reduced by potassium iodide (KI). Rxn (1) S2O82- + 2I- 2SO42- + I2
Rate of reaction was measured by either measuring the disappearance of the reactant or by the appearance of product over time. The researcher decided to use the appearance of iodine. As I2 was produced, S2O32- immediately reacted with it to turn it back to I- Rxn (2) 2S2O32- + I2 S4O6 + 2 I-
Eq (1) Rate = ∆[I2]∆t
This process continued until all of the S2O32- was consumed and converted into tetrathionate. Once all of it was used, I2 started to react with the starch. Rxn (3) I2 + (C6H10O5)n blue complex
It was impossible to measure the change in the concentration of the iodine since its final concentration was not given. Therefore, a substance of known concentration would be of great help in the calculation, and that was S2O32-. The relationship between the [S2O32-] and [I2] can be determined using simple stoichiometry. Eq (2) Δ[I2] = 12(Δ[S2O32-])
Eq (3) Rate = 12(Δ[S2O32-])∆t
Eq (4) ([S2O332-]) = [S2O332-]final – [S2O332-]initial
But since [S2O32-]final was equal to 0,
Eq (5) Rate = 12 – [S2O32-]initial∆t
The rates of reaction for the different runs were also presented in Table 1 above. Sample calculations were provided in sample calculations 2 of Appendix B. If the rates of all the runs and the concentration of the reactants were observed carefully, it was noticeable that as the concentration of the reactants increased, the reaction rate also increased. But the effect was the other way around if their concentration decreased. It happened because as the concentration increased, more molecules were colliding with one another and that resulted in the increase of the frequency of collisions. If there were more collisions, reaction was speed up. The next thing that was needed to be done was finding the order of the reaction. In this experiment, the initial rate method was performed to obtain the reaction orders. Initial rate method makes use of the rate law or rate equation. Formula (3) Rate = k[A]m[B]n
Sample calculations were provided in sample calculation 3 of Appendix B. After the calculation of the reaction orders of each reactants (Peroxodisulfate (VI) ion and iodide ion), the rate law found for the whole reaction was Eq (6) Rate = k [S2O82-][I-]
The proposed mechanism for this reaction that would be consistent with the rate law is I-(aq) + S2O82-(aq) *[I…S2O8]3-(aq) (slow)
*[I…S2O8]3-(aq) + I- I2(aq) + 2 SO42-(aq) (fast)
The first elementary step should be slow because it determined the speed of the reaction. Both of the reactants were also first order, therefore they must both be included in the reactant side of the determining step once.
Now that reaction orders, rate, and reactant concentration were given, the value of rate constant can be found out. This rate constant varies as the temperature changes. The rate constant computed for this experiment under 301 K was 0.06 M-1s-1. Calculation for the rate constant was shown in Sample Calculation 4 of the Appendix B. The unit was M-1s-1 because the whole reaction was in second degree order. Again, since k relies on temperature, the results on this experiment will be invalid if the mixtures have different temperatures.
The next thing done was finding the linear regression of the data. It was used to identify the extent to which there was a linear relationship between a dependent variable and one or set of independent variables. One type of linear regression used was the simple linear regression. It was helpful in predicting the value of the dependent variable using independent variable. In this experiment the concentration of the reactants were the independent variables while the rates of reaction rates were the dependent ones. The equation line of the plot ln of rate and ln of the concentration of the reactants was obtained using the scatter graph in MS Excel. First, the graph of ln rate vs ln [S2O82-] was created and it was shown in Figure 1. The equation of the line was Eq (7) y = 0.937482669x – 8.569597059
It means that for every unit increased in the ln of Peroxodisulfate (VI) ion concentration, the ln of rate increased by .9375. The coefficient of determination (R2) was 0. 3197196. It means that 31.97 percent of the variation in the ln of reaction rate of Peroxodisulfate (VI) ion to iodide ion can be explained by Peroxodisulfate (VI) ion concentration.
Second, the line of the equation of the plot ln rate vs ln [I-] was graphed. It was shown in Figure 2 of Appendix C. The equation of the line obtained was Eq (8) y = 0.873616965x – 9.219000302
It means that for every unit increased in the ln of iodide ion concentration, the ln of rate increased by .8736. The coefficient of determination (R2) was 0.661240957. It means that 66.12 percent of the variation in the ln of reaction rate of Peroxodisulfate (VI) ion to iodide ion can be explained by Iodide ion concentration.
The next factor studied was the effect of temperature to the reaction rate. Raw data was provided in Table 3 of Appendix A. The same thing was done in calculating for the reaction rate as what was done previously. Sample calculation was provided already in sample calculation 4 of Appendix C. The rates obtained for each sets were shown on Table 2. Further information of that part of the experiment was shown in Table 3 of Appendix A.
Run 2| Temperature, K| Reaction Time, s| Rate, M s-1|
Set 1| 301| 93.81| 4.3 x 10-6|
Set 2| 323 | 6.2 | 6.5 x 10-5|
Set 3| 278 | 271.57| 1.5 x 10-6|
Table 2.Effect of Temperature on Reaction Rate
As said previously, rate constant changes as temperature changes. Therefore, different rate constants would be obtained. Sample calculation in finding the rate constant was the same as what was done previously. But this time concentration of run 2 was used. The rate constants obtained were shown in Table 3. Run 2| Temperature, K| Rate constant, M-1s-1|
Set 1| 301| 0.0054|
Set 2| 323| 0.081|
Set 3| 278| 0.0019|
Table 3. Effect of temperature on rate constant
Whenever experiments occur in different temperature, Activation energy (Ea) and Arrhenius constant (A) are always involved. Frequency of colliding molecules cannot be the only factor for the speed up of a reaction. Most molecules just bounce from one another. Therefore, energy is required in order to have a reaction. This is activation energy, Ea, comes in.. Activation energy is the minimum energy required to activate the molecules into a state from which reactant bonds can change into product bonds. Only those with sufficient energy to exceed activation energy can lead to reaction. In order to get Ea, we must get the equation of the line of the plot ln k vs 1/T and by using the Arrhenius equation afterwareds. Formula (4) k = Ae-Ea/ RT
The scatter graph was already presented in Figure 3 of Appendix C. The equation obtained was Eq (5) y = -7357.709686x + 19.89650765.
The value of R2 obtained was 0.908982627. This tells us that 90.90 percent of the variations in ln k can be explained by 1/temperature. The equation obtained could now be used to find Ea. Computation for Ea was already given in sample calculation 5 of Appendix B. The computed activation energy was
Since Ea was already obtained, A was possible to find using also the equation of the line and the Arrhenius equation. Eq (6) ln k = -EaR1T + ln A
Eq (7) y = m x + b
And the equation of the line was
Eq (5) y = -7357.709686x + 19.89650765
Eq (8) ln A = 19.89650765
Eq (9) A = 437,680, 609.4 M-1s-1
Other than frequency of collisions and activation energy, the molecular orientation of the reactant molecules should also be considered and this is what A provides to scientists and researchers.
As observed from the data calculated, as temperature increased, the rate constant increased as well as the reaction rate. The effect of temperature can also be observed in the study entitled “Study on the Reaction Kinetics of Ultrasonic Radiation Non-homogeneous Phase Chitin Deacetylation” done by Zhang C, et al. The ultrasonic radiation provided heat in the deacetylation reaction of the chitin. It was observed that putting such reaction under radiation would increase its reaction rate.  Increasing temperature has caused the average speed of the molecules to increase and therefore their collision frequency also increased. More collision frequency would cause the reaction to speed up
The last part of the experiment was finding the effect of a catalyst in the reaction rate. Rate and rate constant in this part of the experiment were also obtained the way the other rates and rate constants were obtained. The table below shows the results of the calculation made. Run 2| Temeprature, K| Reaction time, s| Rate, M s-1| k, M-s-1| Set 4| 301| 4.73| 8.5 x 10-6| 0.12|
Table 4. Effect of catalyst on reaction rate
Obviously, the catalyst (CuSO4) greatly speed up the reaction rate under room temperature (301 K).
Errors obtained may be caused by the preparation of the reagents. The concentration of the substances especially the reactant ions and thiocyanate could change the rate of reaction. If more thiocyanate was poured on the mixture, it would cause a delay on time because more thiocyanate ions are reacting with the iodine molecules. Slight mistake in recording the temperature from the thermometer could also cause an increase in the experimental reaction rate. Lastly, if the catalyst was not poured as soon as possible, it would also cause in the delay of time before the reaction proceeded.
SUMMARY AND CONCLUSION
To sum it up, an experiment was performed to determine the effects of temperature, concentration of the reactants, and presence of catalyst. The reaction of peroxodisulfate (VI) ion and iodide ion was chosen as the basis of the study. Thiocyanate was used to acquire the initial rates. The experiment was divided into three parts; one for concentration, one for the temperature, and one for the presence of catalyst. The degree of success was 96 percent because there were very small errors obtained from the experiment due to different factors aforementioned. All of the objectives were successfully achieved using the data gathered. The initial rate method was used in determining the rate law of the equation.
The experiment was able to determine the effects of the factors on the rate of reaction. Increasing the concentration of the reactants caused an increase in rate of reaction. Meanwhile, increasing and lowering the temperature of the mixture would speed up and slow down the reaction, respectively. Finally, the presence of catalyst which was cupric sulfate greatly increased the rate of reaction. Base on the data gathered and calculated from the experiment concentration of the reactants, higher temperature, and presence of catalyst would increase the rate of reaction. This study recommends that a statistical tool such as pooled standard deviation be used so as to know how disperse the sets of data are.
 Bayquen, A.V., Exploring Life Through Science Chemistry. Phoenix Publishing House, Inc., 2009, 177.  Mendoza, E.E., Religioso, T.F.,
Chemistry, 3rd edition. Phoenix Publishing House, Inc., 2008, 347.  Bayquen, A.V., Exploring Life Through Science Chemistry. Phoenix Publishing House, Inc., 2009, 178.  Mendoza, E.E., Religioso, T.F., Chemistry, 3rd edition. Phoenix Publishing House, Inc., 2008, 344-346.  Silberberg, M.S., Principles of General Chemistry, 2nd edition. McGraw-Hill, 2010, 514.  Silberberg, M.S., Principles of General Chemistry, 2nd edition. McGraw-Hill, 2010, 527.  Dr.Sample’s Chemistry Classes[Internet].[cited 2013 January 3]. Available from: http://home.comcast.net/~drsamples/ChemWeb401/IodineClockRxnF10.pdf  MITOPENCOURSEWARE [Internet]. Massachusetts: c2005 [cited 2013 January 3]. Available from: http://ocw.mit.edu/high-school/labs/chemistry-intro-to-experimental-chemistry-labs-from-5.302/5_302_iap_2005_3_kinetics_2005b.pdf  Dr.Sample’s Chemistry Classes[Internet].[cited 2013 January 3]. Available from: http://home.comcast.net/~drsamples/ChemWeb401/IodineClockRxnF10.pdf  Bayquen, A.V., Exploring Life Through Science Chemistry. Phoenix Publishing House, Inc., 2009, 181.  MITOPENCOURSEWARE [Internet]. Massachusetts: c2005 [cited 2013 January 3]. Available from: http://ocw.mit.edu/high-school/labs/chemistry-intro-to-experimental-chemistry-labs-from-5.302/5_302_iap_2005_3_kinetics_2005b.pdf  Walpole R.E., Myers, R.H., Probability and Statistics for Engineers and Scientists, 4th edition. Macmillan Publishing Company,1989, 357-358.  Silberberg, M.S., Principles of General Chemistry, 2nd edition. McGraw-Hill, 2010, 529.  Silberberg, M.S., Principles of General Chemistry, 2nd edition. McGraw-Hill, 2010, 530.
 Zhang, C. Cui, Y., Jia, Z., Study on the Reaction Kinetics of Ultrasonic Radiation Non-homogeneous Phase Chitin Deacetylation. Guangzhou,China, 2009. Vol.1.No. 1. 51-53. Available from: http://www.ccsenet.org/journal/index.php/ijc/article/view/418/392.  Silberberg, M.S., Principles of General Chemistry, 2nd edition. McGraw-Hill, 2010, 529.