Introduction

Chemical kinetics involves the study of rates and mechanisms associated with chemical reactions. Reaction rate refers to the change in concentration of a chemical species with time. A reaction mechanism is an exact sequence of steps by which an overall chemical change takes place.

In the laboratory, reaction rate be determined by monitoring the disappearance of a reactant or the appearance of a product over time. In general, the rate of the reaction depends upon the concentrations of reactants. Typically, higher concentration correlates with a higher reaction rate. The exact relationship between rates and concentrations for a reaction is described by the rate law which is a mathematical expression containing the rate, concentrations, and reaction orders.

For the hypothetical equation

aA + bB  cC + dD (1)

in which a, b, c, and d are the stoichiometric coefficients, we can state that

            Rate    [A]m[B]n                    (2)

where m and n are the order of reaction with respect to a given reactant. The rate is directly proportional to and related to the rate by the rate constant k.

            Rate =  k[A]m[B]n                   (3)

The rate law for the reaction under investigation (Equation 1) states that order of reaction with respect to [A] is m and the order of reaction with respect to [B] is n. Note that the exponents m and n are not necessarily related to the stoichiometric coefficients a and b. The reactions orders m and n must be determined experimentally and cannot simply be predicted from an overall balanced reaction.

The rate constant k can be determined for a reaction at a given temperature. The rate constant changes with temperature and can be used for further analysis of a chemical reaction (Week II).

The purpose of this week’s experiment is to determine the overall rate law, including reaction orders and the rate constant for the reaction.

6H+ + IO3– + 8 I– 3 I3– + 3 H2O

This particular reaction is very fast and must be studied indirectly by using an iodine clock reaction to follow the appearance of I3–. Details are given on the experiment website linked below.

Complete the pre-lab worksheet and turn it in to your instructor before the beginning of lab. Late work will not be accepted.

The experiment and background information are located on the following website:

http://web.mst.edu/~gbert/IClock/IClock.html

1. Click on Discussion and carefully read the background information, including the description of the clock reaction. All pertinent reactions are given and explained. Take notes as you read and make sure that you understand how the reactions fit together in this experiment.
2. Click on Experiment.
3. For each experiment, select desired concentrations and then click Mix the Solutions. Then click the Start tab to being timing the reaction. At the moment that the blue color appears, click the Stop tab. The timer must be stopped at EXACTLY the moment that blue appears.

Record this time. Repeat the experiment if you miss the color change. Do at least three trials of each run. When you obtain three times that agree within 5%, average the times and record the average in the data table (Table 1).

4. Run at least four different experiments at 25 °C. Example initial concentrations of each reactant are listed below, but you may use other concentrations if you wish. Run one reaction at a time. Do three trials for each reaction and record the results in the data table included in this document (Table 1).

Run [IO3–]o [I–]o [H+]o Avg. Time (s)*
1 0.005 0.05 2 x 10−5 t1
2 0.010 0.05 2 x 10−5 t2
3 0.005 0.10 2 x 10−5 t3
4 0.005 0.05 4 x 10−5 t4

• time at which blue color appears

1. After data have been gathered for all four runs, calculate the initial rate for reach run. The initial rate, as explained in the background information on the website is given by:

Initial Rate=k[IO_3^- ]^a [I^- ]^b [H^+ ]^c=-1/3 ([H_3 AsO_3])/t

 
Calculations

Calculate the average reaction time for each of the four runs and record on Table 1.

Calculate the reaction rate for each determination using the equation given in Step 5 above. Use your average reaction times for the rate calculations. Record the rates on Table 2 and include correct units with your rates.

Calculate the reaction order with respect to each of the three reactants. Record reaction orders to two digits to the right of the decimal point. You must show your calculations for each reaction order. Record the reaction orders at the top of Table 3. 

Calculate the rate constant k for each of the four determinations using integer reaction orders. Record your results on Table 4. Be sure to show your calculations and include correct units with the rate constants.

On the bottom line of Table 4, write the complete rate law for the reaction you studied in this experiment. Include the average of your calculated values for the rate constant and your reaction orders rounded off to an integer.

Answer the questions on the Post-Lab Worksheet.

Follow all instructions or you will receive a score of zero for this lab.

The attached lab report and worksheets are meant to be printed, filled in, scanned, and returned to your professor. You are welcome to do the calculations on regular paper, but in order to facilitate grading, please draw the tables and reproduce the format of the worksheets.

Obtain a scanning app for your phone. The are many free scanning apps available – Genius Scan is one and the OneDrive app is another – but you can use any app that scans in black and white.

When you are done with your lab assignment, scan it in black and white and save as one pdf. Inspect the document before submitting it to make sure that the format is correct and that all pages are in order. Do not submit photos or other file formats.

 
Kinetics Week I: Lab Report

Table 1. Experimental Data – Reaction Times

Run Temp (°C) [IO3–]o [I–]o [H+]o Time (s)
Trial 1 Time (s)
Trial 2 Time (s)
Trial 3 Time (s)
Average
1
2
3
4

Table 2. Summary of Initial Rate Calculations. Show your work below.

Run [H3AsO3]o Time (s)
Average Temp (°C) Initial Rate
(include units)
1
2
3
4

Make sure to include dimensional analysis and units in all calculations.

Initial Rate Calculation – Run 1
Initial Rate Calculation – Run 2
Initial Rate Calculation – Run 3
Initial Rate Calculation – Run 4

Table 3. Reaction Orders. Show your work bin the table and fill in the blanks. Record reaction orders to two digits to the right of the decimal point.

Runs Used Calculation/Explanation
Order with respect to [IO3–] using Run _ and Run
Calculated Order:
Order with respect to [I–] using Run _ and Run
Calculated Order:
Order with respect to [H+] using Run _ and Run Calculated Order: __

Table 4. Rate Constants and Overall Rate Law. Show your work below.

Run Temp (°C) Rate Constant
(include units)
1
2
3
4
Overall Rate Law =

Make sure to include dimensional analysis and units in all calculations.

Rate Constant Calculation – Run 1
Rate Constant Calculation – Run 2
Rate Constant Calculation – Run 3
Rate Constant Calculation – Run 4

 
Post-Lab Problems

Using the method of initial rates for determining reaction orders, a student finds that:

0.716/0.242=〖1.73〗^x

where x represents the reaction order with respect to a given reactant. Calculate the value of the reaction order x to two decimal places. Show your calculations.

2. For the reaction X + Y + Z  Products, the following data were obtained.
   Hint: Reaction times are inversely proportional to reaction rates.

Run [X]o [Y]o [Z]o Time (s)
1 0.220 0.100 0.224 214
2 0.440 0.102 0.224 107
3 0.439 0.050 0.223 108
4 0.440 0.100 0.447 54

Determine the rate law, including reaction orders with respect to all reactants. Determine the rate constant. In the area provided below, show all your work including units and dimensional analysis. Write your final overall rate law, including the rate constant with units, in the blank given.

Overall Rate Law = ________________