CLARE 102, EXERCISE 5

CHEMICAL KINETICS: RATES OF CHEMICAL REACTIONS

  Introduction

  As described in your lab manual, the rate of any chemical reaction depends on the frequency, intensity, and orientation of particle collisions.

Consider the following set of reactions where the letters (A-L) represent different chemicals.

A   +   B   à   C   +  D                                                                          (1)

E   +   C   à   F  + B                                                                             (2)

E   +   B   +   F   à   G   +   H   +  I                                                       (3)

I   +   C   +   H   à   B   +   J                                                                  (4)

I   +   K   à   L                                                                                       (5)

In order for the end product “L” to be formed, only “A”, “B”, “E”, and “K” must initially be present.  All other chemicals needed for the formation of “L” are produced from reactions between these ‘original’ chemicals and from reactions between the by-products of those reactions.  (To convince yourself of this work backwards from the last reaction).  

However, when “I” is produced in reaction (3), it has two possible fates.  It can react with “C” and “H” or it can react with “K”. Therefore, reactions (4) and (5) are competing. 

If reaction (5) occurs more slowly than reaction (4), then “L” will only be produced in any significant quantity when the supply of either “C” or “H” are exhausted.  Ultimately, the availability of “C” is dependent on the supply of “A” and “B”.

 Experiment Overview

The reactions for today’s experiment are the following:  (Note the letters in bold represent the letters in the generic example previously described.  The reactions are purposely unbalanced for simplicity.)

(A) NaHSO₃  +  (B) H₂SO₄  à  (C) H₂SO₃  +  (D) NaHSO₄                     (1) 

(E) KIO₃  +  (C) H₂SO₃  à  (F) KI  +  (B) H₂SO₄                       (2) 

(E) KIO₃  +  (B) H₂SO₄  +  (F) KI   à  (G) K₂SO₄   +  (H) H₂O  +  (I) I₂                     (3) 

(I) I₂  + (C) H₂SO₃  +  (H) H₂O  à  (B) H₂SO₄  +  (J) HI                     (4) 

(I) I₂  +  (K) starch  à  (L) starch• I₂ complex  (blue-black in color)                     (5) 

Equations (1) through (4) continue until all of the NaHSO₃  is gone.  Then the I₂ is no longer consumed by reaction (4) so I₂ reacts with starch to form a complex that is blue-black in color.

  Today you will mix two solutions containing reactants for this experiment and measure the time until the iodine-starch complex is formed.

Solution 1 is a solution of potassium iodate (KIO₃)

Solution 2 is a solution containing sulfuric acid (H₂SO₄), sodium bisulphate (NaHSO₃), and starch.

(refer to the above set of chemical equations to see how and when these compounds react)

Activity, Part 1

The Effect of Concentration on Reaction Rate

  In order to investigate the effect of concentration, you will prepare dilutions of Solution 1 to vary the concentration of potassium iodate.  The temperature of all of the solutions should be at room temperature.

  General Procedure

1.         Use a pipette to dispense 5.0 mL of Solution 2 into a clean (if you are using a recycled test

            tube, rinse it with RO Water first) test tube. 

2.         Add the appropriate amount of RO Water to the tube using a different pipette.

3.         Add the appropriate amount of Solution 1 to the tube using another pipette. 

4.                  Cap the tube and invert four times.

5.                  Begin timing after the fourth inversion.

6.                  End timing as soon as the solution changes from colorless to blue/black.  (You only need to time to the nearest second.)  Your instructor will have a color standard against which you can compare.

Experimental Procedure. 

The effect of concentration of the potassium iodate can be determined by using different ratios of Solution 1 to RO Water.

Be sure to replicate each treatment THREE times.  Record your data on PAGE 4.

Analysis:  Effect of Concentration

1. Make a generalization concerning the effect of concentration on the rate of reaction.  Give a causal explanation for this generalization?

2. What would happen if you increased the volume of the solution by doubling the amount of all the solutions you put in the tube?  Why?

3.   What would happen if you doubled the amount of Solution 2 without changing either the

      amount of Solution 1 or RO water used?  Why?

4.  Graph the relationship between the concentration of Solution 1 and the reaction time.  To 

     make the graph easier to construct (and interpret), plot the reciprocal of the reaction time

    (1/Average T) rather than the actual average.

5.   Based on the graph and your data for the unknown solution, what is the concentration of this

      solution?

Activity, Part 2   The Effect of Temperature on Reaction Rate

  In order to investigate the effect of temperature you will measure the time of this reaction at a range of temperatures within 20 degrees Celsius of room temperature.  Your instructor will provide you with these solutions.  Use the general procedure outlined in Part I, but use only undiluted solutions (the equivalent of treatment #1 in the previous experiment).  Again make sure the experiment is replicated three times.

 Analysis:  Effect of Temperature

1.         Make a generalization concerning the effect of temperature on the rate of reaction.  Give a causal explanation for this generalization.

2.         Graph the relationship between temperature and reaction time as you did earlier, plotting the reciprocal of reaction time (1/Average T) rather than the actual average.

3.         Based on the graph you constructed and your data for the temperature unknown, what is the temperature of this solution?