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Iodine-Clock Reaction

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Chemical kinetics involving reaction rates and mechanisms is an essential part of our daily life in the modern world. It helps us understand whether particular reactions are favorable and how to save time or prolong time during each reaction. Experiment demonstrated the how concentration, temperature and presence of a catalyst can change the rate of a reaction. 5 runs of dilution and reaction were made to show the effect of concentration on chemical reactions. A certain run from the previous task was twice duplicated to for a “hot and cold” test for reaction rate. The prior run was again duplicated for a test with catalyst. The data obtained was graphed in a linear regression form, using the Arrhenius equation: lnk=-EaR-1T+lnA, reaction rate: rate = k[A]x[B]y and other principles held in chemistry. The graph from the results determined that there is a relation between rate and the previously mentioned factors. The results validated the notion that temperature, concentration and catalysts can manipulate chemical kinetics.


Chemical kinetics is an essential part of our lives. It determines the rates of reactions and how quickly reactants are converted to products. These reaction rates tell us how much time it would take to wait for a change, a physical change [1]. Chemical kinetics can be altered and controlled by five different factors, namely: the temperature at which the reaction occurs, application of pressure, concentration of reactants, nature of reactants and the presence of a catalyst. In this experiment, three of these factors are used to determine and validate their effects to the reaction rate, these factors are the temperature, the catalyst and the concentration of the reactants. S2O82- + 2I-  2SO42- + I2 (1)

The rates of a reaction (2) is defined by the rate law (3), where k is the rate constant, [A] is the concentration of reactant A, [B] is the concentration of reactant B and the powers x and y which gives the order of the reactions of A and B respectively, these values could be computed through the use of initial rates of a reaction. Substituting values from the rate law equation will give us the rate of product formation or decrease in concentration of the reactants in the persulfate-iodide reaction (4). This equation showed how the concentration of reactants modifies the rate of reaction. A + B  C + D (2)

Rate =k[A]x[B]y (3)
Rate = k[S2O82-]x[I-]y (4)
Regarding the effect of temperature on a reaction, we used the Arrhenius equation (5), this where Svante Arrhenius developed the mathematical relationship among activation energy, Ea, absolute temperature and the specific rate constant of a reaction, k, at a certain temperature. In this equation, A is a constant having the same units of k constant and T is the particular temperature as the reaction occurs. It is proportional to the frequency of collisions between reacting molecules that makes up a reaction. R is the universal gas constant expressed with the same units of energy with the activation energy.

In the experiment, 5 different solutions are prepared for their reaction rates and each solution will be called a “run.” The 5 runs will be used to determine the effects of concentration on the reaction. The end of reaction will be determined as I2 reacts with the introduced starch solution, forming a blue solution. Conversely, this reaction is too slow and it needs sodium thiosulfate (Na2S2O3-2) to reduce iodine to iodide before the iodine reaction to starch can take place. Sodium thiosulfate was maintained at a
constant concentration on each run to only show the effects of the concentrations of I- and S2O82- in the reaction [2].

In regards to the temperature-concerned test, two duplicates of the 2nd run in the first set-up was used, conducting the reaction at two highly different temperatures. However, results from the Arrhenius equation above (5) will make it hard to extrapolate the values for graphing, thus a linearized equation (6) was used.

Lnk = -Ea/RT + lnA (6)
The effect of catalyst was tested duplicating the 2nd run again and adding CuS04 as the catalyst. The set-up was at room temperature an the concentrations of the reactants are unchanged to show the change caused by the catalyst METHODOLOGY

Seven solutions were prepared first by class, these are:
1. 500mL 0.2 M KI
2. 500mL 0.2 M KCl
3. 500 mL 0.1M K2S2O8
4. 500 mL 0.1M K2SO4
5. 500 mL 4.0mM Na2S2O3
6. 20 mL of 1% fresh starch solution (0.20g of starch plus 20 mL boiling water, must always be fresh by always heating it) 7. 5 mL CuSO4

Five solutions with the same components but different solutions were be prepared one at a time, each solution was divided into two portions namely, Beaker A and Beaker B. The beakers themselves were filled with the necessary substances and mixtures first before mixing them by pouring Beaker A to Beaker B. The reaction was ensured to be at room temperature. After pouring, the reaction was timed immediately. The timer was stopped after the solution turned to blue.

The concentrations and components of each beaker are presented in this table:
Table 1.

Run 2 was duplicated twice, then one run 2 was placed to a hot bath to reach a temperature of 50 ºC and the other run 2 to an ice bath to reach a temperature of 5 ºC. The contents of the beakers are mixed, and the reaction was timed until the solution turns blue. Catalyst

Another run 2 was made and the set-up where the reaction will take place was ensured to be at room temperature. 4 drops of CuSo4 was added immediately after Beaker A was poured to Beaker B. The reaction was timed until it turned to blue. After the experiment, dispose all the Cu(II) solutions in inorganic waste jars and all the other solutions can be discarded in the sink with copious running water. RESULTS AND DISCUSSION

Table 2. Concentration Effect
| [S2O82-]| [I-]| Time (s)| Rate (M·s-1)| k (M-1·s-1)| 1| 2.00 x10-2| 8.00 x10-2| 52| 1.92 x10-2| 12.0|
2| 2.00 x10-2| 4.00 x10-2| 118.5| 8.43 x10-3| 10.5|
3| 2.00 x10-2| 2.00 x10-2| 227.2| 4.40 x10-3| 11.0|
4| 3.00 x10-2| 4.00 x10-2| 20.5| 1.41 x10-2| 11.75|
5| 4.00 x10-2| 4.00 x10-2| 45.7| 2.19 x10-2| 13.7|

Table 2 shows the concentrations of S2O8-2 and I- and the time of each reaction for each run. Rate of the reaction was calculated by dividing 1M with the time it took to complete the reaction. The constant k was calculated using the rate law (4), the x and y value was determined by the graphs below. Considering most of the k’s calculated by our laboratory class, the supposed real value of k must be 10/Ms for that reaction. No outlier was found, so the calculated standard deviation is 1.22 and a percent error of 17% percent after dividing the absolute value of the difference of the mean (11.79Ms) and 10Ms. These differences must have been obtained from systematic errors such as solution preparation, reaction time when the color changes and the timing of the reaction itself. There could also be random errors from the environment due to the changing of temperature, evaporation of the substances and many more. These results could have been better if the experiment was repeated a few more times.

Figure 1. ln(rate) versus ln[S2082]
Figure 1 shows the relationship between the ln(rate) on the y-axis and ln[S2082] on x-axis of each run. Only runs 2,4 and 5 were considered because the [I-] is constant there. A positive slope was generated after a linear regression was performed and the equation obtained was: y = 7.2838x + 29.82. The generated line after interpolation means that the order of the reactant is 1. This interprets that the relationship between the rate and [S2082] is directly proportional.

Figure 2. ln(rate) versus ln[1-]
Figure 2 shows the relationship between the ln(rate) on the y-axis and ln[1-] on x-axis of each run. Only runs 1,2 and 3 were considered because the [S2082] is constant there. A positive slope was generated after a linear regression was performed and the equation obtained was: y = 0.9367x + 1.2004. The generated line after interpolation means that the order of the reactant is 1. This interprets that the relationship between the rate and [I-] is directly proportional. Overall, the concentration changes on each reactant readily changed the rate of reaction, this validates the notion of concentration being a determining factor in chemical kinetics. To explain why NaS2O32- was added, the reaction between iodine and the starch solutions is very slow, with the presence of S2O32-, a faster reaction was facilitated by reducing iodine to. The reaction between iodine and the starch will produce a blue complex. Table 3. Temperature and Catalyst Effect

Temp. | Set | Temp (K)| Rate (M·s-1)| k (M-1·s-1)|
| 1| 299.15| 8.43 x10-3| 10.5|
| 2| 325.15| 5.12 x10-2| 65.7|
| 3| 278.15| 1.14 x10-3| 1.43|
Catalysis| 4| 299.15| 4.15 x10-1| 595|

Table 3 shows the temperature of sets 1-4 and their respective reaction rates. Sets 1-3 involve the temperature effect, and set 4 is with a catalyst. Note that this shows the direct proportionality of temperature to k. The reaction rates are computed by diving 1M with the reaction time of each set and their k are computed with use of the rate law of the reaction (4). The values of their Ea and A are computed through the equation of the linear regression of sets 1 to 3, this was interpolated with the linearized Arrhenius equation (6). The calculated value of Ea is -8.1646 x 10-4 J/ molK and the Arrhenius constant is 1.238 x 1016, note that the A is always dimensionless. The significance of the activation energy and the Arrhenius constant are the exponential change or the constant miniscule change of the rate, respectively. These will determine the tendency of the reaction, high Ae means a slower reaction and a high A means a slightly faster reaction.

Figure 3. ln(k) versus (1/T)
Figure 3 shows the relationship between the ln(k) on the y-axis and ln(1/T) on x-axis of each set. Only sets 1.2 and 3 were considered because they are the only one to involve drastic temperature change. A negative slope was generated after a linear regression was performed and the equation obtained was: y = 24.4585x + 137.2231. This interprets that the relationship between the k and T is directly proportional. This is because of the collision theory, where the temperature directly relates to the kinetic energy of molecules, as the KE increases more collision happens promoting a faster forward reaction [3]. Note in the catalyst column in Table 3, the rate constant was tremendously increased compared to the rate constant of set 1 who is also at room temperature and has the same percent fractions. This demonstrates the catalyst ‘s ability to increase the rate of a chemical reaction, in this case, CuSO4 was the catalyst that sped up the reaction. The effect of the catalyst works on the activation energies of reaction, the presence of a catalyst lowers the Ae needed for a reaction to proceed [4][5]. CONCLUSION

This reaction mechanism is proposed for the reaction:

I-(aq) + S2O82- (aq) [I…S2O8]3-(aq) (slow)
[I…S2O8]3-(aq) + I-(aq) I2 + 2 SO42-(aq) (fast)

The reaction mechanism satisfied the two requirements: firstly, the sum of the elementary steps gave the overall balanced equation for the reaction;
secondly, the mechanism agreed with the experimentally determined rate law [6] The values of x and y which are solved using the initial rates of the reactions, are both 1, generating the rate law of: rate = k[S2O82-][I-]. The overall order of the reaction is 2, this came from the addition of the rate orders of each reactant. The graphs and tables of the reaction factors in respect to concentration, temperature and presence of catalyst, shows that there is a relationship between these factors and the reaction rate. The experiment was a success in consideration with effect of temperature, concentration and presence of catalyst, it does alter and control the rate of reaction. However, the standard deviation and error percentage in terms of k is relatively large, this may come from the systematic errors like reaction time on the change of color and solution preparation. Random errors could also took place, these could be the evaporation of solutes, contamination and temperature changes. The k values could have been more precise if more replicas of the experiments has been pursued.

[1] Ovelette, Robert. Introduction to General, Organic and Biological Chemistry. 1997. Prentice Hall. p. 236 [2] Wright, Cart. Kinetics of Persulfate Iodide Reaction. 2006. Oxford. Web. December 14, 2012 [3] Whitten K. et. al. Chemistry. 2009. Brooks/Cole, Belmont California. p. 633 [4] Zumdahl,S. Chemical Principles. 1995. D.C Heath and Company, Lexington, Massachusetts p.709-711 [5] Tro, Nivaldo. Introductory Chemistry. 2009. Pearson Education Inc. New Jersey p. 560-563 [6] Zumdahl,S. Chemistry. 2000. Houghton Mifflin Company. Boston Massachusetts. p. 583-586

Rate= 1M/time
Rate= 1M/53 sec
Rate= 1.92 x 10-2 Ms-1

k= rate/([S2O82-][I-])
k= 1.92 x 10-2 Ms-1/(.02M x .08M)
k= 12.0 M-1s-1

y = mx + b
y = 24.4585x + 137.2231
Ea = (m)(R)
Ea = 24.4585 x 8.314 kJ/mol
Ea = -8.1646 x 10-4 J/ molK

Mean= (12 + 10.5 + 11 + 11.75 +13.7)/ 5
Mean= 11.79

SD= { [(12-11.79)2 + (10.5-11.79)2 + (11-11.79)2 + (11.75-11.79)2 + (13.7-11.79)2]/ (4) }0.5 SD= 1.22

% error= (I 11.79-10 I/10) x 100%
% error = 17%


Rate =k[A]x[B]y
Rate = k[S2O82-]x[I-]y
k = A-Ea/RT
Lnk = -Ea/RT + lnA

Mean= (a1 + a2 ….)/ n
SD= {[Σ (x1-mean)2 + (x2- mean)2 ….]/n-1}0.5
% error= (I Exp value-true value I/ true value) x 100%

TABLES: Linear Regression

ln rate vs ln [S2O82-]
Run| rate| y = ln rate| [S2O82-]| x = ln [S2O82-]|
2| 0.00843| -4.776| 0.02| -3.912|
4| 0.014| -4.268| 0.03| -3.506|
5| 0.0219| -3.821| 0.05| -2.995|

ln k vs 1/T

Set| T| x = 1/T| k| y = ln k|
3| 283| 0.00352| 2.475| 0.906|
1| 299| 0.00334| 11.2625| 2.421|
2| 329| 0.00304| 312.5| 5.744|

ln rate vs ln [I-]

Run| rate| y = ln rate| [I-]| x = ln [I-]|
1| 0.019| -3.963| 0.08| -2.525|
2| 0.00843| -4.776| 0.04| -3.218|
3| 0.004| -5.521| 0.02| -3.912|

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