First some safety: The experiment described uses corrosive chemicals and mixtures capable of burning without presence of atmosphere. It should be done outdoors in cotton clothing as fire protection, with chemically resistant gloves, and of course lab goggles. In case of fire, wait for initial reaction to subside, then use a class D extinguisher – do not attempt to put out fires with sand or water.
A recent experiment of mine involved the production of the highly reactive element sodium from sodium hydroxide and another less reactive metal, in this case magnesium. The reaction is as follows:
2Mg + 2NaOH = 2MgO + 2Na + H2
The magnesium (finely powdered) first ignites burning in atmospheric oxygen contained in the reaction vessel. This initial reaction produces enough energy to initiate the primary reaction. The Mg has enough desire for the oxygen contained in the NaOH to actually rip it from the hydroxide group, freeing hydrogen gas and sodium. The reaction is violently exothermic. More than enough energy is released to ignite the hydrogen gas which burns as the reaction progresses. The sodium is produced in a liquid state due to the temperature of reaction, and some of it vaporizes giving the flame a strong orange/yellow color. Despite what is shown in the above video, I have since found that it is better not to grind the hydroxide into a powder. The larger granules slow the reaction, which does not drive off as much of the sodium as a vapor, leaving more to be collected.
To keep the sodium from burning immediately upon production, a loose fitting lid is placed over the reaction vessel, allowing the burning hydrogen to escape, but sealing the vessel off from atmosphere once completed. The lid is then not removed until the reaction cools to air temperature, at which point mineral oil should be added to the resulting slag to prevent further oxidation of the product.
Presumably this experiment could be preformed with other metals acting upon the hydroxide group, such as aluminum. It is also quite likely that this process could free most if not all metals from their hydroxides because the reaction does not involve the metal contained in the hydroxide, but the hydroxide group alone. All hydroxides should therefore be susceptible to separation via magnesium and possibly aluminum.
Ideally, purification and removal of the sodium from the MgO slag would be as simple as heating the mixture under mineral oil to the melting point of sodium, at which point it would gather together. So far I have been unable to do that successfully. It is possible the products needed to be heated longer or they needed to be more finely ground beforehand to free the sodium.