Lab Procedure: Iodine Clock Reaction
Core Concepts
In this lab tutorial, we learn about the iodine clock reaction, including its procedure, underlying chemistry, and data analysis.
Topics Covered in Other Articles
- Lab Safety Rules
- Recrystallization
- Thin Layer Chromatography
- Distillation
- Integrated Rate Laws
What is the Iodine Clock Reaction?
The Iodine Clock Reaction is a classic chemistry experiment that demonstrates many basic principles of kinetics and redox chemistry . For this, the reaction persists as a staple of general chemistry lab demonstrations.
In this experiment, you prepare two simple, transparent solutions. Once the solutions combine, however, the mixture gradually turns from clear to dark blue to near-black. This color change corresponds to the progress of the reaction, which allows you to visually witness the kinetics in a way that most reactions do not provide.
Interestingly, some chemists colloquially call this reaction the “Egyptian Night” experiment. In Egypt, the darkness of nighttime often arrives rather suddenly, similar to rapid dark color change in this reaction.
Iodine Clock Procedure
To perform the Iodine Clock Reaction, you need an iodine salt, a reductant, an oxidant, an acid , starch, and water as a solvent. As mentioned before, these components become allocated between three different solutions according to these specifications:
- First Solution: Starch, Water.
- Second Solution: Iodine Salt, Reductant, Water.
- Third Solution: Oxidant, Acid, Water.
Once the solutions mix, the reaction begins.
The most common variant of the Iodine Clock Reaction uses sodium thiosulfate (Na 2 S 2 O 3 ) as the reductant and hydrogen peroxide (H 2 O 2 ) as the oxidant. Potassium iodide (KI) serves as the salt, while sulfuric acid (H 2 SO 4 ) provides the required acidity. Importantly, gloves, safety goggles, and caution should be observed when using sulfuric acid and hydrogen peroxide to prevent chemical burns.
As we’ll find out in a later section, the kinetics of the reaction depends on the concentrations of acid, iodide, and oxidant. Thus, most lab procedures studying reaction kinetics will vary the concentrations of one or more of these species. Aside from that, the reductant concentration tends to be kept low, as very little is required, while the starch tends to be in excess.
The Chemistry of the Iodine Clock
Iodine clock redox and kinetics.
Before the three solutions mix into one, each ionic species dissociates into their respective ions:
KI → K + + I –
Na 2 S 2 O 3 → 2Na + + S 2 O 3 2-
H 2 SO 4 → H + + HSO 4 –
During the reaction, K + , Na + , and HSO 4 – do not participate, remaining as spectator ions. Once the solutions mix, the hydrogen peroxide oxidizes the iodide into diatomic iodine:
2H + + H 2 O 2 + 2I – → I 2 + 2H 2 O
Importantly, as the reaction produces diatomic iodine, the thiosulfate re-reduces the iodine back to iodide:
2S 2 O 3 2- + I 2 → 2I – + S 4 O 6 2-
This back and forth between iodide and iodine continues until all thiosulfate oxidizes away. Afterward, significant quantities of iodide and iodine exist at the same time. They react with one another to form the triiodide ion:
I 2 + I – → I 3 –
This triiodide ion then forms a complex with the starch. This complex is responsible for the increasing dark blue of the reaction vessel. As a side note, due to the striking dark blue of the complex, a mixture of iodine and iodide called Lugol’s iodine is used to test for trace amounts of starch .
Iodine Clock Kinetics
The first reaction, the oxidation, occurs much slower than the reduction, making it the rate-determining step during that first phase of the reaction. Additionally, once the reduction ceases, the oxidation continues to serve as the rate-determining step, as both the formation of the triiodide and the starch complex occur relatively quickly. Thus, for the entirety of the experiment, oxidation determines the progress of the dark blue hue. This is true even though the starch complex is not directly generated from the oxidation.
Aside from the qualitative observation of the increasingly blue reaction vessel, you can periodically measure the starch concentration through spectrophotometry . The resulting data then allows you to quantify the reaction kinetics.
Kinetic Data Analysis
First, you need to do multiple trials of the Iodine Clock with different concentrations of potassium iodide. Then, you quickly place these samples into a spectrophotometer that records concentrations at consistent time intervals. You’d want to set the spectrophotometer to a frequency similar to 600nm to pick up the dark blue of the starch complex.
Next, you graph your data. You should find that the absorbance of each graph increases linearly with time. This makes sense since chemical reactions always initially proceed at linear rates.
Finally, to determine the reaction order with respect to KI, you take the logarithms of the initial concentrations and reaction rates and generate a log/log graph.
The slope of the resulting slope corresponds to the reaction order in our rate law , due to the properties of logarithms.
RxnRate = k’[KI] n
log(RxnRate) = log(k’[KI] n ) = nlog([KI]) + log(k’)
k’: Relative rate constant (s -1 )
n: Reaction Order of KI
The graph then generates a trendline of y = x – 1.2883, indicating that the Iodine Clock Reaction is first order with respect to KI (n = 1).
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Sodium Thiosulfate Reaction Time with HCl vs. Sulfuric Acid
A PF PlanetHave you tried to find what is the mechanism of the reaction? Or what is its observed kinetics? Answer can be just a google away. FAQ: Sodium Thiosulfate Reaction Time with HCl vs. Sulfuric Acid1. what is the purpose of the sodium thiosulfate reaction time experiment. The purpose of this experiment is to study the rate of reaction of Sodium Thiosulfate with two different acids, HCl and Sulfuric Acid. This will allow us to compare and analyze the effects of different acid concentrations on the reaction rate. 2. What is Sodium Thiosulfate and why is it used in this experiment?Sodium Thiosulfate is a chemical compound with the formula Na 2 S 2 O 3 . It is used in this experiment as a reducing agent and to provide a standard reaction for comparison. It is also used as an indicator for the reaction as it reacts with the acids and causes the solution to turn cloudy. 3. How does the concentration of HCl and Sulfuric Acid affect the reaction rate?The concentration of the acids directly affects the reaction rate. A higher concentration of acid means there are more particles available to react with the Sodium Thiosulfate, leading to a faster reaction rate. Conversely, a lower concentration of acid will result in a slower reaction rate. 4. What factors can influence the reaction time in this experiment?There are several factors that can influence the reaction time in this experiment. These include temperature, concentration of the acids, amount of Sodium Thiosulfate used, and any impurities in the chemicals. It is important to control these factors to ensure accurate and consistent results. 5. How can the reaction rate be determined in this experiment?The reaction rate can be determined by measuring the time it takes for the solution to turn cloudy and recording the results. This can be done by using a stopwatch or a data collection device. The reaction rate can also be calculated by measuring the initial and final concentrations of the reactants and products and using the rate law equation. Similar threads
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 Why is Sodium thiosulphate and sulfuric acid used in the experiment of scattering of light in colloidal solution ?In this experiment, colloidal sulfur is generated by the reaction of sodium thiosulfate and sulphuric acid in a two- step process involving first the formation of thiosulfuric acid followed by its decomposition to sulfurous acid and colloidal sulfur . colloidal suspensions scatter white light strongly. shorter wavelengths are scattered more while longer wavelengths (reds and oranges) are transmitted. scattering is when a molecule absorbs a photon of light, exciting an electron into a higher energy state. when the electron returns to a lower energy state, the molecule emits a photon in a random direction. this randomness produces the scattering known as the tyndall effect. as the sulfur precipitates the higher energy light from the projector will be scattered, making the solution look blue.  The above diagram shows the arrangement for observing scattering of light in colloidal solution, why does the beaker appear blue?  The above diagram is the arrangement for observing scattering of light in colloidal solution, why does the beaker appear orange-red when seen through the slit MN? The scattering of light by colloidal particles present in a colloidal solution is called _______ .  Stack Exchange NetworkStack Exchange network consists of 183 Q&A communities including Stack Overflow , the largest, most trusted online community for developers to learn, share their knowledge, and build their careers. Q&A for work Connect and share knowledge within a single location that is structured and easy to search. Can an iodine clock reaction work without using a strong acid?I'm going to do an iodine clock reaction for a project and we had to submit the materials and safety sheets for the experiment a while ago. I thought I could do one by using hydrogen peroxide, sodium thiosulfate, potassium iodide and starch, but online I'm only seeing experiments that use sulfuric acid or another strong acid of some sort which is worrying me. I talked to my teacher and it's too late to change/add materials, so I'm really hoping that it can work? I'd also like to know what the strong acid does in the reaction, and what would happen (or wouldn't happen) without sulfuric acid? Here's a bit from wikipedia for reference about the reaction: This reaction starts from a solution of hydrogen peroxide with sulfuric acid. To this is added a solution containing potassium iodide, sodium thiosulfate, and starch. There are two reactions occurring in the solution. (note: I didn't try formatting anything so the numbers after the ^ are the charges) In the first, slow reaction, iodine is produced: H2O2 + 2I^− + 2H^+ → I2 + 2H2O In the second, fast reaction, iodine is reconverted to 2 iodide ions by the thiosulfate: 2S2O3^2− + I2 → S4O6^2− + 2I^− After some time the solution always changes color to a very dark blue, almost black. When the solutions are mixed, the second reaction causes the triiodide ion to be consumed much faster than it is generated, and only a small amount of triiodide is present in the dynamic equilibrium. Once the thiosulfate ion has been exhausted, this reaction stops and the blue colour caused by the triiodide – starch complex appears.
 You wrote the answer yourself: the hydrogen peroxide consumes hydrogen ion according to the reaction $\ce{H2O2 + 2 I^- + 2\color{blue}{H^+} -> I2 + 2 H2O}$ You need the acid to provide the hydrogen ions, especially since if you allow the solution to become basic then the iodine disproportionates (see "Chemistry and Compounds" section, "Iodine Oxides and Oxoacids" subsection). In terms of pure chemistry, maybe you could have used a moderately strong acid such as phosphoric acid or a bisulfate salt. But then the limited dissociation of such species might make the iodine formation reaction slower and thus slow down the "clock". Looking at the Wikipedia article on this reaction, one alternative is to use a peroxydisulfate salt instead of hydrogen peroxide as your oxidant. This reacts with iodide ion according to $\ce{S2O8^{2-} + 2I^- -> 2 SO4^{2-} +I2}$ with thiosulfate again as the reducing agent to regenerate iodide ion. Unlike hydrogen peroxide or the other oxidants, the above reaction does not consume hydrogen ions and is your best bet to run the reaction without a strong acid addition. However, peroxydisulfates have their own hazards (one of which is requiring a strong sulfuric acid solution to make it in the first place), and if you make the common choice of potassium peroxydisulfate rather than the ammonium salt you could have solubility issues. 
Your AnswerReminder: Answers generated by artificial intelligence tools are not allowed on Chemistry Stack Exchange. Learn more Sign up or log inPost as a guest. Required, but never shown By clicking “Post Your Answer”, you agree to our terms of service and acknowledge you have read our privacy policy . Not the answer you're looking for? Browse other questions tagged acid-base reaction-mechanism experimental-chemistry equilibrium redox or ask your own question .
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Catalysis of a sodium thiosulfate and iron(III) nitrate reactionIn association with Nuffield Foundation
Investigate the effect of transition metal catalysts on the reaction between iron(III) nitrate and sodium thiosulfate In this experiment, students compare the rate of reaction between iron(III) nitrate solution and sodium thiosulfate solution when different transition metal ions are used as catalysts. The catalysts used are copper(II), cobalt(II) and iron(II) ions. Iron(III) ions are reduced to iron(II) ions in the presence of sodium thiosulfate. The reaction proceeds via a dark violet unstable complex but gives a colourless solution with time. Students can do this experiment in pairs or small groups. If each pair of students attempts this experiment, large volumes of both the iron(III) nitrate solution and the sodium thiosulfate solution will be required.
Access to 0.1 M solutions of the following (see note 8 below):
Health, safety and technical notes
 Source: Royal Society of Chemistry As the solution changes from dark violet to colourless, the cross underneath the measuring cylinder will become visible
Teaching notesInitially the iron(III) and thiosulfate ions form an unstable complex (which is dark violet in colour): Fe 3+ (aq) + 2S 2 O 3 2– (aq)→ [Fe(S 2 O 3 ) 2 (H 2 O) 2 ] – (aq) Over time the complex is consumed as thiosulfate (acting as a reducing agent) reduces iron(III) to iron(II) ions. Transition metal ions can catalyse this reduction process at different rates. If too much catalyst is used then the reaction proceeds instantaneously. It is important that students only use one drop of catalyst. It is possible to set up this experiment using a light sensor and data logging. The data logging software should show the colour change occurring on a graph. This gives more information than the standard end point approach. The rate of change can be measured from the slope of the graph or the time taken for the reaction to occur. Student questionsSome possible questions to ask students include:
Additional informationThis is a resource from the Practical Chemistry project , developed by the Nuffield Foundation and the Royal Society of Chemistry. This collection of over 200 practical activities demonstrates a wide range of chemical concepts and processes. Each activity contains comprehensive information for teachers and technicians, including full technical notes and step-by-step procedures. Practical Chemistry activities accompany Practical Physics and Practical Biology . © Nuffield Foundation and the Royal Society of Chemistry
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Related articles 4 ways to teach redox in terms of electrons2024-07-03T05:06:00Z By Kristy Turner Use these teacher-tested approaches to help learners gain a deeper understanding of redox reactions  New method grows larger diamonds2024-05-31T08:24:00Z By Nina Notman Use this real-world context when teaching about giant covalent structures  Mastering titration apparatus2024-05-07T08:38:00Z By Kristy Turner Use this poster, fact sheet and classroom activity to show learners the names and uses of equipment they’ll encounter in this practical No comments yetOnly registered users can comment on this article., more experiments.  ‘Gold’ coins on a microscale | 14–16 yearsBy Dorothy Warren and Sandrine Bouchelkia Practical experiment where learners produce ‘gold’ coins by electroplating a copper coin with zinc, includes follow-up worksheet  Practical potions microscale | 11–14 yearsBy Kirsty Patterson Observe chemical changes in this microscale experiment with a spooky twist.  Antibacterial properties of the halogens | 14–18 yearsBy Kristy Turner Use this practical to investigate how solutions of the halogens inhibit the growth of bacteria and which is most effective
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Procedure. Put 50 cm 3 of sodium thiosulfate solution in a flask. Measure 5 cm 3 of dilute hydrochloric acid in a small measuring cylinder. Add the acid to the flask and immediately start the clock. Swirl the flask to mix the solutions and place it on a piece of paper marked with a cross. Look down at the cross from above.
ObjectiveRate of Reaction of Sodium Thiosulfate and Sulphuric acidKinetics Study on the Reaction between Sodium Thiosulphate and Sulphuric acidThis video is ...
2. action of Sodium Thiosulfate and Hydrochloric Acid continuedDiscussionSodium thiosulfate react. ion 1).Na2S2O3(aq) + 2HCl(aq) → S(s) + SO2(g) + 2NaCl(aq) Equation 1The kinetics of the reaction can be analyzed by graphing the. oncentration of Na2S2O3 as a function of both reaction time and 1/time. A plot of concentration versus time gives a ...
Third Solution: Oxidant, Acid, Water. Once the solutions mix, the reaction begins. The most common variant of the Iodine Clock Reaction uses sodium thiosulfate (Na 2 S 2 O 3) as the reductant and hydrogen peroxide (H 2 O 2) as the oxidant. Potassium iodide (KI) serves as the salt, while sulfuric acid (H 2 SO 4) provides the required acidity ...
The reaction is: Na 2 S 2 O 3 (aq) + 4H 2 O 2 (aq) → Na 2 SO 4 (aq) + H 2 SO 4 (aq) + 3H 2 O (l) The sulfuric acid produced by the reaction neutralises the sodium hydroxide (buffered by the sodium ethanoate) and gives the observed colour changes. If the reaction is done with 20 volume hydrogen peroxide, the reaction is slower than with the ...
2. Add 10 mL of 3-M sulfuric acid to each flask. 3. Use a pipette to place exactly 2.00 mL of liquid bleach into one of the flasks, and stir by swirling the solution. Fill a 50-mL burette with standard sodium thiosulfate solution, and use the thiosulfate solution to titrate the triiodide formed during the reaction of bleach with iodide.
Add 4.5 ml sulfuric acid and stir the solution until the potassium iodate dissolves. Dilute to 1 L. ... add about 10 g sodium thiosulfate to neutralize the iodine to iodide. Stir until the mixture becomes colorless. ... making it suitable for home school experiments. Blue bottle reaction: This redox reaction changes from blue to clear. While ...
In this demonstration, colloidal sulfur is generated by the reaction of sodium thiosulfate and hydrochloric acid in a two- step process involving first the formation of thiosulfuric acid followed by its decomposition to sulfurous acid and colloidal sulfur(3). 2 H+ (aq) + S2O32- (aq) --> H2S2O3 (aq) H2S2O3 (aq) --> H2SO3 (aq) + colloidal sulfur
Put 10 cm 3 of sodium thiosulfate solution and 40 cm 3 of water into a conical flask. Measure 5 cm 3 of dilute hydrochloric acid in a small measuring cylinder. Warm the thiosulfate solution in the flask if necessary to bring it to the required temperature. The object is to repeat the experiment five times with temperatures in the range 15-55 °C.
Revision notes on Required Practical: Investigating the Effect of Concentration on Rate of Reaction for the AQA GCSE Chemistry syllabus, written by the Chemistry experts at Save My Exams.
PP041 - The thiosulfate-acid reaction: rate and concentration. Small scale 'disappearing cross' methods for investigating the effect of concentration on the rate of the sodium thiosulfate - acid reaction. Includes use of an alkaline 'stop-bath' to neutralise used reaction mixtures which helps to minimise fumes.
Sep 1, 2017. Sodium. In summary, the reaction time between sodium thiosulfate and HCl is faster compared to sulfuric acid. This is due to the higher reactivity of HCl and its ability to dissociate into ions more easily, resulting in a more rapid reaction. However, sulfuric acid can still be used as an alternative and its slower reaction time ...
An alternative protocol uses a solution of iodate ion (for instance potassium iodate) to which an acidified solution (again with sulfuric acid) of sodium bisulfite is added. [3]In this protocol, iodide ion is generated by the following slow reaction between the iodate and bisulfite: IO − 3 + 3 HSO − 3 → I − + 3 HSO − 4. This first step is the rate determining step.
In this experiment, colloidal sulfur is generated by the reaction of sodium thiosulfate and sulphuric acid in a two- step process involving first the formation of thiosulfuric acid followed by its decomposition to sulfurous acid and colloidal sulfur . Colloidal suspensions scatter white light strongly.
This experiment will allow students to find out some interesting chemical reactions of sodium thiosulphate, record, observe, and understand this compound. Students will induce reactions between sodium thiosulfate and other chemicals. This practical takes place in three parts, with each part showing learners a new side of this complex substance.
In the first, slow reaction, iodine is produced: H2O2 + 2I^− + 2H^+ → I2 + 2H2O. In the second, fast reaction, iodine is reconverted to 2 iodide ions by the thiosulfate: 2S2O3^2− + I2 → S4O6^2− + 2I^−. After some time the solution always changes color to a very dark blue, almost black. When the solutions are mixed, the second ...
It neutralises any remaining acid and the sulfur dioxide reacts with the water to produce sulfuric acid. If the indicator is showing the acidic colour, refresh the stop bath by adding more sodium carbonate solution. Ensure the room is well ventilated. A microscale version of the experiment is available from CLEAPSS.
Abstract figure legend Sodium thiosulfate (STS) rescues the pronephros phenotype of pdx1 morphants through compensatory upregulation of nitric oxide (NO) metabolism. ... primarily related to energy metabolism such as oxidative phosphorylation and citric acid cycle, were likewise downregulated (Fig. 4A). The experiment consisted of five clutches ...
Reaction of sulfuric acid and magnesium ribbon. Repeat steps 1-3 of the first experiment, using sulfuric acid in place of sodium hydroxide solution. Add one 3 cm piece of magnesium ribbon. Stir with the thermometer and record the maximum or minimum temperature reached.
Place the cylinder back on top of the cross. Using a 50 cm 3 measuring cylinder, measure 50 cm 3 of iron (III) nitrate solution. Pour the iron (III) nitrate solution into the sodium thiosulfate solution, and start the timer. An immediate dark violet solution is observed which turns colourless after a few minutes.