This collection of classroom resources features all the pupil activities from our Developing expertise in teaching redox chemistry course for teachers. This collection is most valuable to those who have attended this course and wish to put into practice with their students some of the ideas and activities presented as part of that event.
Examine the reactions between various metals and metal salt solutions
Illustrate the displacement of copper from copper(II) sulfate solution using aluminium foil.
This experiment compares the colours of three halogens in aqueous solution and in a non-polar solvent. These halogens also react to a small extent with water, forming acidic solutions with bleaching properties.
Transform methylthioninium chloride from blue to colourless and back again by mixing it with glucose and shaking the solution, then letting it settle.
Illustrate the idea of competition reactions by reacting copper(II) oxide and zinc metal together, causing an exothermic reaction.
A teaching resource on Chemical Change, supported by video clips from the Royal Institution Christmas Lectures® 2012.
Determine the formula of copper(II) oxide by reducing it using hydrogen or methane.
Treat your students to a spectacular demonstration as aluminium and iodine are catalysed by water.
Filter paper soaked in potassium iodide solution which also contains starch and phenolphthalein is placed on an aluminium sheet which forms one electrode of an electric circuit. The other electrode is used as a ‘pen nib’ to ‘write’ on the filter paper. When this electrode is made positive and the ...
An interesting introduction to the electrolysis of brine (sodium chloride solution). Students use Universal Indicator to help them follow what is happening during the reaction.
This demonstration shows that an ionic salt will conduct electricity when molten but not when solid. Zinc chloride is used - this will melt at Bunsen burner temperatures.
This activity looks at rusting in the context of shipwrecks. It has different demands to the traditional experiment to show the factors needed for rusting to occur.
Introduce your students to the idea that different oxidation states of transition metal ions often have different colours and that electrode potentials can be used to predict the course of the redox reactions.
Illustrate the charging and discharging of a lead-acid cell to show the relationship between the electrical energy put into the cell and the energy released.
A simple demonstration of catalysis also introducing the idea of an activated complex and to allow discussion of the mechanism of catalysis.
This activity introduces oxidation numbers by giving a conceptual foundation for them in terms of electron accounting and polar bonds. It then shows how the model used so far needs refining.
Many reactions between gases and solids are suitable for demonstrations and class practicals. Making reaction tubes is an excellent lesson in physical chemistry in its own right as well as being cheaper than buying in expensive material.
Demonstrate a clear increase in mass as iron wool is heated in air on a simple ‘see-saw’ balance.
Observe an endothermic reaction as solid hydrated barium hydroxide is mixed with solid ammonium chloride to produce a liquid that evolves into ammonia gas. The temperature drops dramatically to about -20 °C.
Produce a small explosion in your classroom by electrolysing water then re-combining the hydrogen and oxygen gas.
The background and chemistry of burning calcium with hydrogen and oxygen to make ‘limelight’.
This experiment involves the reaction of a metal with the oxide of another metal. When reactions like these occur, the two metals compete for the oxygen. The more reactive metal finishes up with the oxygen (as a metal oxide). If the more reactive metal starts as the oxide then no ...
Students reduce iron(III) oxide with carbon on a match head to produce iron in this small scale example of metal extraction. The experiment can be used to highlight aspects of the reactivity series.
Prove that aluminium is a more reactive metal than iron by demonstrating the highly exothermic reaction between aluminium and iron(III) oxide resulting in molten iron.
Students heat copper(II) oxide in a glass tube while passing methane over it. The copper(II) oxide is reduced to copper. If the reactants and products are weighed carefully the formula of the copper oxide can be deduced. This could also be used simply as an example of reduction.
A mixture of alcohol and air in a large polycarbonate bottle is ignited. The resulting rapid combustion reaction, often accompanied by a dramatic ‘whoosh’ sound and flames, demonstrates the large amount of energy released in the combustion of alcohols.
In this demonstration experiment, a mixture of glycerol (propane-1,2,3-triol) and potassium manganate(VII) crystals bursts into flame, giving off clouds of steam, after a short time lag.
This is a demonstration that shows the reactions of Group 1 metals in air and in chlorine. It does not clearly show the trends in reactivity of Group 1 metals, which are better demonstrated by the reactions in water, which follow on well from this demonstration.
Gases give rise to particular hazards so great care must be taken when preparing, collecting or testing.