Alkali metal roulette

Students generally make three requests after seeing the reaction of alkali metals with water: ‘Put in a bigger piece!’, ‘Can we see what happens with caesium?’ and ‘What happens if you do it in acid?’.

Safety prohibits the use of bigger pieces, while cost and practical considerations prohibit the use of caesium. But, perhaps surprisingly, the reaction with concentrated hydrochloric acid can be safely performed.

Kit

• Potassium (release flammable gases in contact with water)
• Sodium (release flammable gases in contact with water)
• Concentrated hydrochloric acid (corrosive)
• 100 cm³ of propan-2-ol (highly flammable and irritant to eyes) 
• Two 125 mm x 16 mm borosilicate test tubes
• Retort stands, bosses and clamps
• Knife to cut metals
• Paper towel
• Tweezers
• Two safety screens

Preparation

Use safety screens to protect both the audience and the demonstrator. Clamp the test tube and load it with concentrated hydrochloric acid to a depth of 4 cm, which should make it about one third full. Ensure the acid is fresh: bottles opened some time ago may have evaporated off hydrogen chloride, reducing the concentration and paradoxically increasing the risk associated with this demonstration. The bore of the tube must be no larger than 16 mm to ensure the concentration of hydrogen gas is such that it will not explode inside the tube. You may wish to set up a camera focused on the top of the liquids to help the audience see the changes in the metal and the formation of the precipitate.

In front of the class

Ensure the audience members are all positioned at least 3 metres away from the demonstration and wearing eye protection. Cut off a piece of metal about the size of a peppercorn, pat off any oil with a piece of paper towel and place the metal into the tube using the tweezers. Spectators are likely to predict a dramatic reaction, but instead a precipitate of metal chloride will rapidly form, protecting the metal from good contact with the acid. The metal fizzes gently and glows red hot while the evolved hydrogen ignites at the mouth of the tube and burns with a lilac flame – if it does not ignite spontaneously, you can use a splint.

Teaching goal

Any chunk of alkali metal in a classroom environment is likely to be mostly covered in a layer of metal oxide. Even if it is well cleaned, more oxide, as well as metal hydroxides, will form on contact with air and moisture. The reactions of the alkali metals with water are heterogenous and limited by the contact the underlying metal has with water. As such, the presence of a product film will further inhibit the reaction so the solubility of these metal oxides and hydroxides is crucial to determining how quickly the metal can react.

Although it’s unclear to what extent film formation may be affecting the reactions of metals in the open laboratory, experiments have shown it takes place, at least under low pressure with gaseous water, and it is likely that the same can be said in liquid water, if only for lithium.¹

This demonstration shows how product solubility can produce unexpected results; the reaction of concentrated HCl being less dramatic than ethanoic acid as the metal ends up sitting on a layer of salt that precipitates out of solution and slows the reaction with the bulk acid beneath.

The table shows why this reaction should only be attempted with concentrated HCl, which results in the reaction proceeding in a controlled manner as the metal sits on a protective layer of precipitated sodium chloride. For the salts of other acids, things don’t work out quite as nicely. The ethanoate salt of potassium is even more soluble than the hydroxide, leading to an extremely vigorous reaction. The solubility of the nitrate salt of potassium is low, but the oxidising effect of the nitric acid and the higher hydrogen ion concentration (15.7 mol dm^-3,compared with 11.7 mol dm^-3  for HCl) mean that this should not be attempted either. The low solubility of the sulfate salts looks tempting but at 18 mol dm^-3  for sulfuric acid, and it being diprotic as well as oxidising, this would also be a dangerous choice.

Solubility is also likely to play a role in the differing reactivities of the alkali metals with water. Product solubilities vary wildly to the point that, at least for lithium, the solubility of surface impurities is likely to slow its reaction further when compared with the other metals. Once solvated, products must then diffuse away from the metal surface. Lower down the group, the ions are more mobile because their ability to attract water molecules around themselves is less, and thus the hydrated ion radius is smaller. 

Both of these solubility effects enhance the reaction of the later alkali metals with water.

This demonstration was originally developed by Colin Chambers and published in a supplement to Ted Lister’s Classic Chemistry Demonstrations. A special risk assessment was produced by Peter Borrows and Bob Worley.³