Iodine Clock Reaction
🔒 LAB RESOURCES:
Student Procedure [PDF] [DOC]
Teacher Annotated Procedure [PDF]
Complete Lab Guide [HERE]
GoogleSheet Data & Analysis [HERE]
SUMMARY/ OVERVIEW:
Students determine the rate law and rate constant for the reaction between iodide and persulfate ions. The "clock" effect occurs because a small amount of thiosulfate instantly consumes iodine ($\text{I}_2$) until the thiosulfate is used up. Once $\text{I}_2$ builds up, it reacts with starch, causing a sudden, measurable blue-black color change.
ESTIMATE TIME ⏰: 40-60 minutes
Running Trials (40 minutes)
SAFETY PRECAUTIONS:
Introduction
Chemical kinetics is the branch of chemistry that is concerned with the mechanisms and rates of chemical reactions.
The mechanism of a chemical reaction is a description of what happens to each molecule at a very detailed level—which bonds are broken, which new bonds are formed, and how the three-dimensional shapes of the chemicals change during the course of the reaction. The rate of the reaction is a measure of its speed. The rate of a chemical reaction can be measured by how quickly the reactants disappear, or by how quickly the products are generated.
The iodine clock reaction is a favorite demonstration in chemistry classes because it has an element of drama. Two clear solutions are mixed, producing a new clear solution. Then, after a period of several seconds, the solution turns dark blue. A demonstration of this reaction is shown in the video below.
Experiment Reactions
As mentioned, chemical kinetics measures how fast a reaction is occurring. For most chemical reactions, the rate is so fast that special equipment is needed to measure it. For the iodine clock reaction, on the other hand, the rate can be easily measured by monitoring the color change of the reaction. To perform the iodine clock reaction, you will mix potassium iodide, sulfuric acid, starch, and thiosulfate. The time it takes for the reaction mix to turn blue will be measured with a timer.
The reactions that form the basis for the iodine clock reaction are shown below.
The "Clock" Mechanism
Step 1: The Primary Reaction (The "Timer")
The overall reaction being studied is the relatively slow oxidation of iodide ions ($\text{I}^-$) by persulfate ions ($\text{S}_2\text{O}_8^{2-})$:
$$\text{2I}^-(\text{aq}) + \text{S}_2\text{O}_8^{2-}(\text{aq}) \rightarrow \text{I}_2(\text{aq}) + 2\text{SO}_4^{2-}(\text{aq})$$
This reaction is the "timer." Its rate depends on the initial concentrations of the reactants, $[\text{I}^-]$ and $[\text{S}_2\text{O}_8^{2-}]$. During this period, $\text{I}_2$ is produced at a constant, slow rate.
$$\text{Rate} = k[\text{I}^-]^x[\text{S}_2\text{O}_8^{2-}]^y$$
Step 2: The Secondary Reaction (The "Stopper")
To make the timing measurable, a small, fixed, and limiting amount of thiosulfate ions ($\text{S}_2\text{O}_3^{2-}$) is added to the mixture. Thiosulfate reacts instantaneously and preferentially with any $\text{I}_2$ produced in Step 1, immediately consuming it:
$$\text{I}_2(\text{aq}) + 2\text{S}_2\text{O}_3^{2-}(\text{aq}) \rightarrow 2\text{I}^-(\text{aq}) + \text{S}_4\text{O}_6^{2-}(\text{aq})$$
Result: As long as there is $\text{S}_2\text{O}_3^{2-}$ present, the concentration of free $\text{I}_2$ remains essentially zero, and the solution remains colorless or pale yellow. The thiosulfate acts as a scavenger, preventing the reaction from moving to the visual end point.
Step 3: The End Point (The "Alarm")
The solution also contains a starch indicator. Starch reacts specifically with free $\text{I}_2$ to form a large, intensely colored complex.
When the slow reaction (Step 1) has produced exactly enough $\text{I}_2$ to consume all of the limiting $\text{S}_2\text{O}_3^{2-}$, the thiosulfate is completely used up.
Once the $\text{S}_2\text{O}_3^{2-}$ is gone, the $\text{I}_2$ produced by the slow reaction finally begins to build up in concentration.
This free $\text{I}_2$ instantly reacts with the starch, causing a sharp, rapid change to an intense blue-black color.
The measured time interval ($\Delta t$) from mixing to the color change is the precise time needed for the slow reaction (Step 1) to generate the fixed amount of $\text{I}_2$ required to neutralize the initial, known amount of thiosulfate.