CHEMICAL KINETICS: THE IODINE CLOCK REACTION

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ABSTRACT

This paper is a study to form an idea of how certain experimental factors affect the rates of reactions. The experiment conducted demonstrated how the temperature, the concentration and the addition of a catalyst affect the rate of a reaction. Through the preparation of different runs with varying concentrations of reactants with a constant concentration of S2O32- and the use of a starch indicator as a basis for the end of the reaction, the chemical kinetics behind the reaction of S2O82- and I were observed.

The different runs were timed to determine the rate constant, the reaction rate, the order of the reaction and the rate law. From the results obtained it was concluded that the reaction followed the second order rate law, with both the sodium thiosulfate and iodide having first order. The rate of reaction was also found to be directly affected by temperature; a higher temperature resulted in a faster reaction rate, and the addition of a catalyst increased the rate of reaction.

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INTRODUCTION

Chemistry deals with the prediction and the manipulation of chemical reactions, and the rate of the reaction plays a significant role in this practice. There are reactions that can happen in less than a second, while there are others that take a long period of time.

Chemical kinetics is the branch of chemistry that is concerned with the mechanisms and rates of chemical reactions. The description of what happens to each molecule on a detailed level is the mechanism of a chemical reaction. The measurement of a reactions speed is known as the rate of reaction. The rate of a chemical reaction is measured by how fast the reactants disappear, or by how fast the products are created [1].

The study of the rate of a chemical reaction occurs and the factors that affect this rate. Reactions depend on the chemical nature of the reactants and the rates of each reaction are unique given a set of conditions. There are four factors that can influence the rate of a reaction one is the the concentration and physical states of the reactants, another is the temperature of the system and the other is the presence of a catalyst. As discussed by the collision theory, molecules of the reactants should collide and mix in order to react.

The rate law is a mathematical equation that identifies the reactions progress. The equation relates the rate of the reaction to the rate constant (k) and the concentration of the reactants raised to their reaction orders.

The Iodine Clock Reaction Experiment aims to determine how the concentration, the temperature and the addition of a catalyst effects the rates of reaction using the rates of different runs of an Iodine and persulfate reaction mixture.

The experiment includes the preparation of the following solutions in Table 1.

Table 1. The concentration of the reagents used

Reagent Concentration

Potassium Iodide (KI)0.2M

Potassium Chloride (KCl)0.2M

Potassium Persulfate (K2S2O8)0.1M

Potassium Sulfate (K2SO4)0.1M

Sodium Thiosulfate (Na2S2O3)4.0mM

Copper (II) Sulfate (CuSO4)0.1M

Starch Solution1% w/v

The concentrations of the reactants were varied for each run as presented in table 2. The varied amount of the solutions were done in order to determine the effect of the concentration of the reactants to the rate of reaction.

Table 2. The volumes of the reactants for the different runs

RUNBEAKER ABEAKER B

KI(mL)KCl(mL)K2S2O8(mL)K2SO4(mL)Na2S2O3(mL)

1100555

255555

32.57.5555

4557.52.55

5551005

For Set 1 all runs were prepared. Eight drops of the starch solution were added to beaker B and the solutions in beaker A were poured into beaker B marking the start of the reaction and the recording of the reaction. The starch solution mixed with the Iodine to form the blue complex indicating that the Iodine was formed and marking the end of the reaction.

For set 2 the observation of the effect of temperature on the rate of reaction was tested. Two sets of Run 2 (in table 1) were set up. This time the environment of the solutions were in different temperatures. One set of the second run was placed in an ice bath while the other set was heated; the set-ups had to reach the temperature of -4 and 50 respectively. The addition of the starch and the mixing of the solutions indicated the start of the reaction, and the blue complex formed marked the end.

This experiment dabbled in the effect of a catalyst to the rate of reaction. A catalyst lowers the activation energy of the reaction. The activation energy is the minimum amount of energy required for the reaction to take place [2]. To observe the effect of a catalyst 1M of the catalyst, CuSO4, was added for a third set of Run 2 was prepared. The reaction was recorded as the catalyst was added and the contents of Beaker A and B were mixed, and the blue complex that formed marked the end of the reaction.

RESULTS AND DISCUSSION

The Iodine Clock reaction records how long it takes for a fixed amount of the thiosulfate ions to be consumed. In the reaction there are two processes simultaneously occurring. The first process is a slow reaction that produces the iodine (Equation 1).

S2O82- (aq) + 2I- (aq) I2 (aq) + 2SO42- (1)

The Iodine produced is never seen because it is immediately used up in a very fast process that reduces it back into the colourless iodide ion (Equation 2).

I2 (aq) + 2S2O32- (aq) S4O62- (aq) + 2I- (aq) (2)

The Iodine is formed slowly until the thiosulfate ion is used up. When the color of the solution changes into a blue complex it is the result of the complete consumption of the thiosulfate ion (Equation 3). The moment that the solution turned into a blue complex the timer was stopped.

I2 (aq) + starch blue complex (3)Reaching the blue complex point of the reaction depends on the concentration of the reactants and the rates of the two reactions. If there is anything that will speed up the first reaction then the overall reaction will also speed up. The rate of reaction would not be significant if the blue complex did not appear or if the reaction turn ed blue as soon as the reactants were mixed. The thiosulfate ions regulate the rate at which would react to the starch solution, therefore the rate of reaction depends on it [3].

In order to observe the end of the reaction the concentration of the thiosulfate must be lower than the concentration of the persulfate and iodide ions. Other wise, if the thiosulfate had a greater concentration the persulfate would be consumed first and the change in color would not be visible [4].

The reaction is dependent on the thiosulfate ion, therefore the rate of the reaction may be computed using the rate of the consumption of the thiosulfate (Equation 4).

rate = -[S2O32-] (4) 2t

The negative sign in the equation is used because the thiosulfate ion is used up and the concentration will have a negative change giving the rate a positive value. The time is doubled because of the coeffecient of thiosulfate in Equation 2. [5]

Although, the rate of the reactions were determined not using the rate law. Instead using the reciprocal of the time elapsed when the blue mixture was formed because it would produce the same amount of I2 over the time and is directly proportional to the rate of reaction.

Table 3. Effect of the Concentration on the rate of reaction

S2O82-MI-MS2O32-MRxn time, sRate, Ms-1

10.0330.21.33x10-327.833.5X10-2

20.0330.21.33x10-344.632.2x10-2

30.0330.051.33x10-3101.249.9x10-3

40.0330.11.33x10-325.084.0x10-2

50.0330.11.33x10-318.835.3x10-2

In set 1, the concentration of persulfate was constant for runs 1, 2 and 3 while the concentration of iodide was the constant for runs 2, 4 and 5. The final concentration of the ions were calculated from the molar concentration and volumes of each solution in every run. Using the initial rate method the reaction orders of persulfate and iodide were calculated.

The reaction order calculated for persulfate was close to 1 and for iodide was 1 making the overall reaction order to be 2. The rate law is expressed as the rate constant multiplied by the concentration of the reactants and raised to their reaction orders. [6] Which looks like this

Rate = k [S2O82-][I-](5)

Reaction orders may also be determine graphically through plotting of ln (rate) vs. ln (concentration of the reactants). Figure 1 is the graph of the obtained equation y = 0.318x + 1.443 and linearity value of R2 = 0.9966. The determined value of the slope is the reaction order of S2O82- to be 1.

Figure 1. The relationship of [S2O82-] and the rateln rate

Figure 2 is the graph obtained from the equation y = 2.189x + 1.964 with a linearity value of R2 = 0.9952. The slope of the line is the reaction order of I- which is approximately equal to 1.

Figure 2. The relationship of [I-] and the rate

ln rate

Based on the values obtained from the graphs the rate law can be established as

Rate = k [S2O82-]1[I-]1

SUMMARY AND CONCLUSIONS

The Iodine Clock Reaction was an experiment that tested three factors that affect the rate of reaction, more specifically: the concentration, the temperature and the use of a catalyst. Varying the concentration of the reactants may affect the average reaction time as well as the rate of the reaction. Increasing the temperature increases reaction rates because of the disproportionately large increase in the number of high frequency collisions. It is only these collisions if they possess the activation energy for the reaction that will result in a reaction. The addition of a catalyst lowers the activation energy of the reaction allowing it to proceed faster.

The values that are obtained may vary depending on the factors that may affect the reaction. For one, an inaccurate timer may be a liable source for error, other factors could be the thermometer which could change the temperature that is recorded. The results obtained and presented through this experiment are not too far off from other experimental findings. The differences may be accounted for as errors and misreadings of data.

REFERENCES

[1] Chang, R. 2007. Chemistry: Ninth Edition. New York: McGraw Hill. 549-553.

[2] "CHEM 1112 . Kinetics of the Persulfate iodide Clock Reaction. Web. .

[3] [4] Harris, D.C. 2011. Quantitative Chemical Analysis: Eight Edition. New York: W.H. Freeman Company. 347.

[5] [6] Upadhyay, S.K. 2006. Chemical Kinetics and Reaction Dynamics. New York: Anamaya Publishers. 2-6

APPENDIX A - List of Tables

Table 1. The concentration of the reagents used

Table 2. The volumes of the reactants for the different runs

Table 3. Effect of the concentration on the rate of reaction

APPENDIX B - List of Figures

Figure 1. The relationship of [S2O82-] and the rate

APPENDIX B- SAMPLE CALCULATIONS

Rate Orders

Rate = k [S2O82-]m[I-]n

3.5X10-2 = k(0.033)m(0.2)n n = 0.7955 2.2x10-2 k(0.033)m(0.1)n

4.0X10-2 = k(0.033)m(0.1)1 m = 1 2.2x10-2 k(0.033)m(0.1)1

Rate Constant

Run 2: 0.033 M persulfate 0.1 M iodide

k = rate = 2.2x10-2 k = 6.6666667 [I-][S2O82-] (0.033)(0.1)

Solution Preparations

For 500mL 0.2M KI

500mL x 1L x 0.2M x 166g = 16.6g 1000mL 1mol

REFERENCES