Fluorine (F)
Fluorine (F) is a highly reactive and electronegative element that belongs to Group 17 of the periodic table, known as the halogens. With an atomic number of 9, it is the most reactive element, readily combining with almost all substances except for noble gases like helium, neon, and argon. At room temperature, fluorine is a pale yellow gas with a pungent odor and is notably dense compared to air. The element's remarkable reactivity necessitates careful handling, as it can react explosively with water to form hydrofluoric acid, a potent corrosive substance.
Fluorine was first isolated in 1886 by French chemist Henri Moissan, whose method of electrolysis remains in use today. Fluorine is commonly found in nature not in its free form but as part of minerals such as fluorite and fluorapatite, with significant production in China and Mexico. In everyday applications, fluorine compounds are integral to dental health, as fluoride is added to toothpaste and drinking water to prevent cavities. Additionally, fluorine plays a crucial role in the pharmaceutical industry, enhancing the efficacy and selectivity of many drugs. Despite its benefits, certain fluorinated compounds, such as chlorofluorocarbons (CFCs), have raised environmental concerns, leading to international efforts to regulate their use due to their detrimental effects on the ozone layer.
Subject Terms
Fluorine (F)
- Element Symbol: F
- Atomic Number: 9
- Atomic Mass: 18.9984
- Group # in Periodic Table: 17
- Group Name: Halogens
- Period in Periodic Table: 2
- Block of Periodic Table: p-block
- Discovered by: Ferdinand Frederic Henri Moissan (1886)
Fluorine is one of the six halogen elements that constitute Group 17 of the periodic table. The other elements in this group are chlorine, bromine, iodine, astatine, and element 117 (its temporary name is ununseptium, which is the word for one-one-seven in Latin). Fluorine, which is represented by the symbol F, has an atomic number of 9. Of all the elements in the periodic table, it is the most reactive and electronegative element. (Electronegativity is the propensity of an atom to attract a pair of bonding electrons.) At room temperature, fluorine is a pale yellow gas that reacts with almost all organic and inorganic substances.
The history of fluorine begins in 1529 when Georgius Agricola, a German physicist, described fluorspar, a fluorite-bearing mineral. But it wasn’t until 1886 that the French chemist Henri Moissan finally managed to isolate fluorine. He was able to obtain elemental fluorine by electrolyzing a solution of potassium hydrogen fluoride and hydrogen fluoride. In 1906, Henri Moissan won the Nobel Prize in Chemistry for his achievement. Shortly after returning from Stockholm, where he received his award, Moissan suddenly passed away. However, his method of isolating elemental fluorine continues to be used to this day.
At the time of the Second World War, uranium hexafluoride was being used to separate uranium isotopes, and several organic fluorine compounds of industrial use were developed that helped fluorine achieve significance as an industrial chemical. In everyday life, compounds of fluorine are used in toothpaste and drinking water. Teflon, a household name, is polytetrafluoroethylene—that is, a synthetic fluoropolymer of tetrafluoroethylene. Food cooked in Teflon-coated frying pans does not stick to the surface. Hydrofluoric acid is used in the etching of light bulbs and other glassware.
Physical Properties
At room temperature, Fluorine is a pale yellow gas with a pungent odor. At 1.695 grams per liter, it is approximately 1.3 times as dense as air. Fluorine is a gas in its standard state at 298 K. Fluorine gas turns into liquid at –188.13°C, and the liquid form of the element turns into a solid at –219.61°C. The specific heat of an element is the amount of energy required to raise the temperature by one degree. For fluorine at 298 K, the specific heat is 0.82 joules per kilogram. Some specimens of the mineral fluorite (CaF2) exhibit fluorescence when exposed to ultraviolet light.
Chemical Properties
With three exceptions, fluorine reacts with most of the elements in the periodic table. The three exceptions are helium, neon, and argon. Extreme caution is necessary when handling fluorine because it reacts explosively with most compounds, especially water, to produce the corrosive solution hydrofluoric acid. Chemical substances fail to free fluorine from its compounds, and for this reason, scientists have found it extremely difficult to isolate it.
Although fluorine has eighteen known isotopes, only one of them is stable. Since fluorine-19 is the only stable isotope of fluorine, the element is considered to be a monoisotopic element. Fluorine-18 is the longest-lived isotope of fluorine, with a half-life of about 110 minutes. The long half-life of fluorine-18 allows it to serve as an important source of positrons. Fluorine-15 is the least stable of the fluorine isotopes, with a half-life of 4.1 × 10-22 seconds. Fluorine-18 decays by positron emission 97% of the time and by electron capture 3% of the time. In both cases, it yields stable oxygen-18.
Applications
In nature, fluorine rarely occurs in its free state. Fluorite, fluorapatite, and cryolite are the most common fluorine minerals. Fluorine occurs abundantly in Earth’s crust; it is almost as abundant as barium or manganese. Roughly constituting 0.06% of Earth’s crust, fluorine is the thirteenth most common element. For all commercial purposes, fluorite (CaF2) is generally used as a source of fluorine. China and Mexico have historically produced the majority of the world’s fluorite. In the United States, its main source is Illinois, but some of the other states where fluorite is mined include Kentucky, Ohio, Tennessee, New Mexico, New York, and New Hampshire.
Fluorine is the twenty-fourth most abundant element in the universe, and it is created when a neutrino collides with a neon atom during a supernova. It is also the product of another cosmic event: when a blue Wolf-Rayet star of more than forty solar masses experiences a stellar wind that blows the fluorine out of the star sooner than hydrogen or helium can destroy the fluorine. Almost all of the fluorine in the human body occurs as the fluoride that shows up in the calcified tissues of bones and teeth.
Commercially, electrolysis using a solution of potassium hydrogen difluoride in anhydrous hydrofluoric acid yields fluorine. Because it does not conduct electricity, pure hydrofluoric acid cannot be used to produce fluorine.
Fluorides are used to prevent dental cavities. The most commonly used fluoride is sodium fluoride. In order to prevent tooth decay, fluoride is added to drinking water. As much as 95% of the fluoride that is added to public water supplies is obtained from phosphorite rock. Tooth enamel is made up of hydroxyapatite. Fluoride replaces the hydroxyl molecule, thereby making teeth resistant to bacterial attack. One of the strongest inorganic acids, hydrofluoric acid, is used both as a cleaning agent and also in the etching of glass. Hydrofluoric acid is produced when fluorite is treated with sulfuric acid.
Fluorine is widely used in the field of medicine. Because the properties of molecules show a significant change when fluorine atoms are added to the preexisting molecules, drugs that include fluorine become more selective, show an increase in efficacy, and are easier to administer. Among all the drugs available in the market today, around one-fifth of them contain at least one fluorine substituent. (A substituent is an atom or a group of atoms that substitutes for another atom or group in a molecule.)
Fluorite is employed as a flux, or cleaning agent, that is used in smelting. It is also used in the production of certain glasses and enamels.
Chlorofluorocarbons, or CFCs, were used as refrigerants before they were shown to produce harmful effects that damaged the ozone layer of the atmosphere. Because of these environmental effects, the Montreal Protocol that was adopted in Montreal on September16, 1987 mandated a phasing out of the use of CFCs.
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