Catalysis of a reaction between sodium thiosulfate and iron(III) nitrate solutions
The rate of reaction between
iron(III) nitrate solution and sodium thiosulfate solution is compared 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.
Stopclock or timer
Dropping pipette. Use the type of teat pipette usually fitted to Universal indicator bottles that does not allow squirting
Glass measuring cylinder (100 cm
Measuring cylinder (50 cm
Access to 0.1 M solutions of the following (
Cobalt(II) chloride solution, (TOXIC), drops
Copper(II) sulfate solution, drops
Iron(II) sulfate solution, drops
Iron(III) nitrate solution, 250 cm 3
Sodium thiosulfate solution, 250 cm 3
Refer to Health & Safety and Technical notes section below for additional information.
Health & Safety and Technical notes
Read our standard health & safety guidance
Wear eye protection.
Cobalt(II) chloride solution, CoCl 2(aq), (TOXIC) - see CLEAPSS Hazcard and Recipe Book.
Copper(II) sulfate solution, CuSO 4(aq) - see CLEAPSS Hazcard and Recipe Book.
Iron(II) sulfate, FeSO 4(aq) - see CLEAPSS Hazcard and Recipe Book.
Iron(III) nitrate solution, Fe(NO 3) 3(aq) - see CLEAPSS Hazcard and Recipe Book. If iron(III) nitrate is not available, iron(III) chloride, 0.1 M , or iron(III) ammonium sulphate, 0.1 M, can be used instead.
Sodium thiosulfate solution, Na 2S 2O 3(aq) - see CLEAPSS Hazcard and Recipe Book.
It is important that the concentrations of the solutions are accurate. If higher concentrations are used the experiment proceeds too quickly. It is useful if each group of students has access to their own supply of solutions, this prevents contaminating the bulk supply. The catalyst solutions can be provided in dropping bottles and the iron(III) nitrate and sodium thiosulfate solutions in 500 cm
a Draw a cross on a piece of scrap paper and put it underneath the 100 cm 3 measuring cylinder so it can be seen when looking down the cylinder from the top.
b Using the 100 cm 3 measuring cylinder, measure 50 cm 3 of sodium thiosulfate solution. Place the cylinder back on top of the cross.
c Using a 50 cm 3 measuring cylinder, measure 50 cm 3 of iron(III) nitrate solution.
d 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.
e Look through the reaction mixture from above until the cross can first be seen. Stop the timer and record the time.
f Repeat this experiment, but add one drop of catalyst to the iron(III) nitrate solution before mixing. Test the various catalysts in the same way.
g Record the times for no catalyst and all the catalysts tested.
Initially the iron(III) and thiosulfate ions form an unstable complex (which is dark-violet in colour):
3+(aq) + 2S 2O 3 2-(aq) → [Fe(S 2O 3) 2(H 2O) 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.
Here are some possible questions to ask students.
1 Which is the best catalyst?
2 Why were only very dilute solutions of the catalysts used?
3 Could you slow the reaction down? If so, how?
Health & Safety checked, 2016
This Practical Chemistry resource was developed by the Nuffield Foundation and the Royal Society of Chemistry.
© Nuffield Foundation and the Royal Society of Chemistry
Page last updated October 2015