Demonstrations designed to capture the student's imagination
This activity, involving an oxygen cyclinder, is unlikely to be covered by model risk assessments generally adopted in UK schools. A special risk assessment, obtained directly from an organisation such as CLEAPSS or SSERC, would be required.
A simple and safe class demonstration of the reduction of iron ores (Fe2O3, Fe3O4) to iron involves burning a match until the end is charcoalised, dipping the burnt end into water and then into some iron oxide. The iron oxide-coated match head is then introduced into a roaring Bunsen flame where the reduction takes place, producing enough iron for the cooled match head to be attracted to a magnet. Although a good example of what happens in a blast furnace, this demonstration lacks because the students don't see and appreciate what I would call 'pure iron'.
Here I use a gentle supply of oxygen at the surface of the reaction mixture to generate the high temperatures (>1000ºC) necessary for the direct reduction of iron oxide with carbon.1 The resulting sparks, smell and bright light make for a dramatic demon-stration, and the resulting lump of iron can be handled when cooled.
- Steel or ceramic crucible supported in a sand bath;
- 1m Stainless steel/glass (10mm diameter) tube attached to an oxygen cylinder fitted with a flash back arrester (pressure required <1bar (100kPa));
- Iron oxide, Fe2O3; wood charcoal;
- Bunsen burner, heat-resistant mats and tongs, tinted safety goggles, lab coat and protective gauntlets.
Three quarters fill the crucible with a 50:50 mix of iron oxide and wood charcoal. (Note: activated charcoal is useless because it is too fine and blows away when the oxygen stream is introduced.) Compact the mix by gently pressing with your finger, and embed the crucible in the sand bath. Heat the top of the crucible with a roaring Bunsen flame until the charcoal starts to glow red. Turn off the Bunsen and put on the tinted goggles. Using the stainless steel tube, which should be at least 1m long for safety, direct a gentle stream of oxygen gas at ca 0.5bar (50kPa) pressure over the surface of the mixture. Too much pressure and you'll blow the mixture away, too little and the heat produced will not be sufficient. When the reaction has finished you will see a lump of molten iron in the crucible. Care must be taken not to burn this iron by continuing to supply oxygen. After cooling, use tongs to tip out the lump of iron onto a heat-resistant mat.
The sparks generated by the reaction travel some height above the crucible so never stand directly above it. Tinted safety goggles must be worn to protect your eyes from the intense light generated by the reaction. Students must stand well back (3m) and wear safety goggles. Although less bright than burning magnesium ribbon (no uv/blue light is given off), students must be told not to stare directly at the flame. Place heat-resistant mats around the sand bath to protect the floor. You must ensure there are no oxygen leaks from the tubing connecting the cylinder to the stainless steel tube and that care is taken to separate the oxygen from any source of ignition.
Instead of a stainless steel tube, a glass tube can be used. Tinted safety goggles are available from school lab suppliers for ca £12 and from most DIY stores. Care must be taken when judging the oxygen flow rate. I judge the correct flow rate by blowing the oxygen on my cheek until I can feel a gentle breeze.
When teaching the extraction of metals I start with native metals such as gold, and then move up the reactivity series. Silver oxide can be heated easily to produce silver. Copper is more problematic but lengthy roasting of a mixture of copper oxide and charcoal will do the job, or more simply thermal decomposition of the oxide directly. Iron, being more reactive requires, more effort, which this demonstration shows well.
I also link the thermal decomposition of iron oxide to some simple entropy calculations with my Y13 students. By making some assumptions, they can predict the temperature of thermal decomposition of iron oxide:
Fe2O3(s) → 2Fe(s) + 3/2O2(g)
ΔrH° = +824kJmol-12
ΔS°surroundings = -ΔH/T
ΔS°system = ΔS°products - ΔS°reactants = ((27.3 × 2) + (102.5 × 3))Jmol-1K-1 - 87.4Jmol-1K-1 = +274.7Jmol-1K-1
So for the reaction to be feasible ΔStotal must be greater than zero;
ΔStotal = ΔSsystem + ΔSsurroundings
0 = 274.7Jmol-1K-1 + (-824kJmol-1/T)
1. D. F. Shriver, P. W. Atkins and C. H. Langford, Inorganic chemistry, p230. Oxford: OUP, 1990.
2. R. D. Harrison, Revised Nuffield advanced science book of data. Harlow: Longman, 1984.