If chemistry is about reaction, then fluorine stands out as its star performer. Nothing enters into chemical reaction with more vigor than elemental fluorine. Chemists now declare that fluorine forms compounds with every single element except neon and helium, the list extending all the way up to elements number 106 (Seaborgium). This list includes the heavier noble gases (those that resist reaction with anything else), and even the members of its own group – the halogens.
The fact that fluorine is so reactive means that most of the time you are dealing with the compounds of fluorine, and these stand out as distinct in many ways. To prise fluorine out its compounds is a Herculean feat, and the history of its discovery rivals that of the expedition to the South Pole. Sir Humphrey Davy had a scent of something special, as he writes, “it appears reasonable to conclude that there exists in the fluoric compounds a peculiar substance, possessed of strong attractions for metallic bodies and hydrogen”. But elements number 9 remained a gap in Mendeleev’s periodic table until the year 1886, when French chemist Henri Moissan finally succeeded in isolating the gas, using electrolysis and coating his instruments with a highly resistant platinum-iridium alloy. It was the culmination of a string of laboratory accidents that left many deaths in its wake, described as the “fluorine martyrs”.
The atomic structure of fluorine explains everything. Like all the halogens, it contains five electrons in its outer p shell. This means that the addition of only one further electron delivers the crucial landmark of a completed p shell (6 electrons), the stable configuration. Therefore, fluorine is said to have an extremely high electron affinity (328 kJ per mole), which means that it will vigorously grasp at any electron available in order to form a compound. This is the case with all the halogens, but nothing matches fluorine, which gains its extra advantage from the smallness of the size of its atom. This means that the positive nucleus exercises a stronger pull than those of other halogens, whose successive shells of electrons insulate the nucleus.
The propensity of elements to capture electrons and thereby form negative ions is called its electronegativity. Fluorine’s is stands at the highest at 4. By the same token, fluorine forms highly stable ionic compounds. In normal conditions it is a pale yellowish gas in its diatomic molecular form. In other words, two fluorine atoms join covalently to form the fluorine molecule, F2. Naturally occurring fluorine is mono-isotropic, i.e. is found with only one isotope, of atomic weight 19. Sir Humphrey Davy not only anticipated its existence, but also named it, from its ore – fluorite (calcium fluoride).
The compounds of fluorine find widespread use in industry. In many of these instances the usefulness stems from the strong ionic bonds by which fluorine attaches to its compounds. One of most well-known is as “fluoride” in toothpaste to prevent tooth decay. The fluoride ions present in toothpaste react with the enamel to form fluorapatite, lending to the teeth a highly resistant outer layer. Fluorine is also added to drinking water for the same purpose. In the same vein, Teflon (poly-tetra-fluoro-ethylene) lends a non-stick layer to cooking utensils.
In fact, fluorine products are used in the same way in which the fluorite ore was in metallurgy as far back as the 16th century. In smelting the metal from its ore, the addition of fluorite added fluidity to the mix (hence the name, from the Latin fluer, “to flow”). At the same time it protected the metal from atmospheric corrosion through the formation of a protective layer of its fluoride. Today, artificial cryolite (sodium hexafluoaluminate) is used in the production of aluminium, which accounts for a significant part of the modern fluorine industry.
Fluorine products have found many uses in medicine. A significant proportion of pharmaceutical products now contain fluorine, which is introduced in order to boost reactivity in the bloodstream. Many anesthetic gases, such as isoflurane and sevoflurane, are fluorine derivatives of hydrocarbons. Nuclear magnetic resonance (NMR) scans, used to obtain highly accurate diagnostic images, uses the mono-isotropy of natural fluorine along with its high magnetic moment. Another radioactive isotope of fluorine is used in PET (positron emission tomography) scans.
Fluorine was also used was in the Manhattan Project, carried out in the University of Chicago during the Second World War in order to develop the atom bomb. In order to obtain the “fissionable” uranium 235 from the more abundant uranium 238, its fluorine derivative uranium hexafluoride was used. The light weight of fluorine made sure that there was insufficient mass gradient between the two fluorides, which allowed for purification of fissionable uranium through repeated diffusion.
On a more controversial note is the use of products known as chlorofluorocarbons (CFCs) in refrigeration and aerosols. The company DuPont launched the fluorine industry after it discovered the remarkable refrigerant properties of freon (chloro-fluoro-ethane) in 1941. Substantial use over the following decades meant that these products ended up in the upper atmosphere where they begin to affect the ozone layer. In this way the extreme reactivity of fluorine ends up afflicting great damage to the planet. The ozone layer, which protects the Earth from deadly ultraviolet rays, is crucial to life. Its depletion through CFCs triggered enormous environmental campaigns, culminating in the historic Montreal Protocol on 1990 which totally banned their worldwide production.
Extremely activity also means that fluorine products are generally poisonous to life. The presence of fluorine in 100 ppm is considered hazardous, and four times that as fatal. The toxicity derives from fluorine is strong affinity towards calcium in the bloodstream. Pesticides and rat poison also contain fluorine.
