Iodine Clock Reaction – Notes for Instructors
Ministry Expectations and Application to Energy Changes and Rates of Reaction Unit
This lab is designed to address a student’s understanding of rates of reaction as outlined in the curriculum documents. More specifically, it falls under section “D3. demonstrate an understanding of energy changes and rates of reaction”. This experiment involves students diluting standard solutions and monitoring how the rate changes over time with consecutive dilutions. Specifically, this lab demonstrates how concentration can affect the rate of a reaction. Students use the data they collect during the lab to determine the order of the reaction based on their results.
The iodine clock reaction involves mixing two colorless solutions together and then monitoring the mixture for a colour change. The reactions occurring in this experiment are:
IO3- + HSO3- à 3SO42- + I- + 3H+ (slow)
6H+ + IO3- + 5I- à 3I2 + 3H2O (slow)
I2 + starch à blue black complex (fast)
In this protocol, iodide ion is generated by the slow reaction between the iodate and bisulfite. Since this is the slowest step of the reaction, it is also the rate determining step. The iodate in excess will oxidize the iodide generated forming iodine. The iodine is immediately then reacts with the starch to form a blue black coloured solution. The reaction is sensitive to dissolved oxygen and hence reactions using stock solutions should be carried out fairly soon after they have been prepared. If the distilled water that is used is degassed, the solutions will last longer.
Safety Considerations for the Instructor
The preparation for this lab can be done on a bench top. None of the materials are volatile or produce fumes. H2SO4 is a strong acid and gloves and goggles should be worn when handling this chemical. Note only a very small amount is required in this lab. Should any contact occur with the skin or eyes rinse with copious amounts of water and consult a doctor as soon as possible.
Safety Considerations for the Students
Students should wear goggles and protective clothing such as a lab coat when performing this experiment. Ensure that no students drink any of the solutions. Should a student come in contact with any of the solutions prepared for this experiment, either on their skin or in their eyes, rinse the affected area with copious amounts of water and go to a hospital immediately.
Preparing the Solutions
NOTE: Water should be degassed Ar or N2. Hence it is best to pour about 200 or 300 ml into a large beaker so that the 4 litre bottle is tightly capped. This will avoid dissolving oxygen into the stock water.
Chemicals Required:
- KIO3
- Distilled Water (degassed with N2 or Ar if possible)
- Soluble starch
- NaS2O5
- H2SO4
Equipment Required:
- 100 ml beakers (x4)
- 500 ml beaker (x2)
- Glass stirring rod
- Hot plate
- 100 ml graduated cylinder
- 100 ml volumetric flasks (x2 with Teflon/plastic stoppers)
- Digital scale
- Tape, marker
Procedure for Preparing A: 0.020 mol/L potassium iodate solution
- weigh 0.43 g of KIO3 into a beaker (100 ml is a good size)
- use a glass stirring rod to pulverize the powder in the beaker after
- add about 30 ml of water to the beaker and dissolve the solid
- pour this solution into one of the 100 ml flasks and dilute to 100 ml
- label the flask as “Solution A”
Procedure for Preparing B: 0.00100 mol/L sodium bisulfite
- in a 100 ml beaker, boil about 40 ml of water and then add to the boiling water 0.40 g of soluble starch
- stir the mixture and try to dissolve as much of the solid as possible
- if it clumps up, that is ok; stop stirring after 2 minute of boiling
- pour the solution into a 100 ml graduated cylinder and dilute with water up to the 90 ml mark
- the starch solution should be translucent and it will probably also be cloudy; that is ok
- in a separate 100 ml beaker, weigh 0.02 g of NaS2O5 and 0.05 g of H2SO4 (one drop is fine) and add about 5 ml of water
- pour this solution into the graduated cylinder and then pour the contents of the cylinder into a 100 ml volumetric flask; dilute to100 ml
- label the flask as “Solution B”
Before the Lab
Students should be familiar with rate laws and what 0, 1st and 2nd order reaction diagrams look like. As well, students should be familiar with dilution calculations. Before the lab, it is advisable to review quickly the relationship between reaction rate and concentration and also about rate laws and how they are developed.
Running the Lab
- Run the lab with students in pairs
- A sample class room set-up is shown below with benches/tables setup and labeled 1 through 6
- Depending on the size of the tables/benches, one can have 4 students (2 pairs) at each bench
- Place one 500 ml beaker at each bench for waste collection and a 200 ml beaker with clean, degassed distilled water with a labeled dropper (label = clean water)
- Each pair of student gets:
o 1 x test plate
o 1 x 100 ml beaker of solution A
o 1 x 100 ml beaker of solution B
o 1 x 100 ml beaker for clean water to rinse the pipettes
o a stop watch
o 2 x droppers
o a clean sheet of white paper
- layout the equipment for the lab: group each item together i.e. beakers, test plates, etc. and have the students come up to the front lab bench individually by student bench to get their materials
- have the students label two of the 100 ml beaker “A” and “B” and then go around the desks and provide them with the required chemicals from stock solutions that were pre-prepared
- ensure students also have the required work sheet to go with the lab
Answers to Post-Lab Questions
1. Assume that each drop has a volume of 0.2 ml. Given that the concentration of Solution A (KIO3) is 0.02 M, calculate the initial concentration of IO3- in each of the wells at the moment both solutions are completely mixed together. Put your answers in tabular form.
Use C1V1=C2V2. The solutions are more dilute as more water is added, so the concentration should go down. Students may make the mistake of not taking V2 = 20 ml.
# of drops of water added / Initial Concentration0 / 0.020 M
2 / C1 = 0.02 M
V1 = 0.2 ml/drop x 8 drops = 1.6 ml IO3-
C2 = ?
V2 = 20 ml
C1V1/V2 = C2
(0.020 mol/L)(1.6 ml)/20 ml = 0.0016
4 / C1 = 0.02 M
V1 = 0.2 ml/drop x 6 drops = 1.2 ml IO3-
C2 = ?
V2 = 20 ml
C1V1/V2 = C2
(0.020 mol/L)(1.2 ml)/20 ml = 0.0012
6 / C1 = 0.02 M
V1 = 0.2 ml/drop x 4 drops = 0.8 ml IO3-
C2 = ?
V2 = 20 ml
C1V1/V2 = C2
(0.020 mol/L)(0.8 ml)/20 ml = 0.0008
8 / C1 = 0.02 M
V1 = 0.2 ml/drop x 2 drops = 0.4 ml IO3-
C2 = ?
V2 = 20 ml
C1V1/V2 = C2
(0.020 mol/L)(0.4 ml)/20 ml = 0.0004
2. What is the general trend in the reaction time as less of solution A is used?
The reaction times get longer. If they do not, then the solutions may be oxidizing. But the student should offer some explanations as to why they are not getting longer reaction times.
3. Use the graph paper on the back of this sheet to plot the initial concentrations you calculated from Question 1 versus the reciprocal of the reaction time you obtained from your experiment.
Students should use a proper scale and should use as much of the graph paper as possible. Graph should be linear. Students should then realize that if the graph is linear, the order of the reaction must be equal to 1.
4. Based on your graph, what is the order of the reaction?
Reaction should be first order with respect to IO3-.
Other Resources:
This reaction can also be done using peroxide. Here is a very good website that could be used as a guide should you wish to go through the peroxide route instead of using sulfuric acid.
http://chempaths.chemeddl.org/services/chempaths/?q=book/General%20Chemistry%20Textbook/Chemical%20Kinetics/1526/rate-equation