Fluorine
Elementary fluorine probably is the strongest oxidizer known. It is an exceedingly poisonous and exceedingly corrosive very pale yellow gas, which eats virtually everything, including such inert things as glass and hard plastics. Elementary fluorine even is capable of attacking the inert gasses xenon and krypton and it can oxidize elementary oxygen to a positive oxidation state!
Fluorine can only be prepared and handled with great difficulty and there are only few labs in the world, which have the equipment and expertise to do research, involving elementary fluorine. The home chemist can only dream of this.
Compounds of fluorine frequently are remarkably inert, especially when a completely fluorinated species is involved. Examples of this are NF3, CF4, SF6. Higher strongly fluorinated carbon-fluorine compounds also are remarkably inert, teflon being a well-known nice example.
On the other hand, ionic fluorides and hydrofluoric acid are quite dangerous and very reactive. Despite of the fact that hydrofluoric acid is a weak acid (it is only weakly ionized in aqueous solution), it is among the most dangerous mineral acids. It produces exceedingly painful wounds and even dilute solutions must be handled with great respect.
Fluorides are moderately interesting from the point of view of the home chemist. Fluorine has no aqueous redox chemistry, but it is interesting as a ligand in many coordination complexes.
Hydrofluoric acid
should not be present in the average home lab. The benefits of using this
compound really are not worth the risks of storing and handling this acid.
The following less dangerous fluorine compounds are available for the public, but still, they are quite toxic (except calcium fluoride, due to its insolubility) and on acidification also become insidiously corrosive:
- sodium fluoride, NaF
- calcium fluoride, CaF2
- potassium fluoride, KF
- potassium bifluoride, KHF2
- ammonium bifluoride, NH4 HF2
All of these are white solids. Of these, calcium fluoride, is not interesting at all. It is very inert and only concentrated hot acids are capable of slowly dissolving this, with the liberation of gaseous hydrogen fluoride. Sodium fluoride sparingly dissolves in water (approximately 3 grams per 100 ml). Potassium fluoride dissolves in water more easily. The bifluorides are compounds with the HF2– anion. Although less dangerous than hydrofluoric acid, these still are very dangerous and one should be really careful not to get in contact with solutions of bifluorides.
Calcium fluoride is available at many pottery and ceramics suppliers. The other fluorides are available as special order at some raw chemical photography suppliers. Ammonium bifluoride might also be available as etchant in certain well-equipped art stores.
Solutions of fluorides and bifluorides slowly attack glass! If one experiments with these, then the experiments should be conducted in cheap transparent plastic beakers, otherwise your glasswork will be ruined. Especially if some acid is added, the etching effect is quite noticeable.
Summarizing: Fluorides can add something to the home lab, but they must be handled with care and one must know the risks of using them.