Fluorine

Fluorine is a highly reactive chemical element with the symbol F and atomic number 9.

It is the lightest member of the halogen group. At standard conditions, it exists as a pale yellow diatomic gas. This means it is made of two atoms bonded together. Fluorine is famous for being extremely reactive. It will react with almost all other elements. The only exceptions are the light noble gases. Because of this intense reactivity, elemental fluorine is highly toxic. It is a powerful substance that plays many roles in modern industry and science.

The way fluorine behaves is driven by its unique atomic structure. A fluorine atom has nine electrons. Two electrons fill the inner shell, leaving seven in the outer shell. This outer shell requires only one more electron to become full.

Because of this, fluorine has a very high electronegativity. This term describes how strongly an atom pulls on electrons. Fluorine has the highest electronegativity of any reactive element. It also has a high electron affinity, which is the tendency to capture an electron. This drive to complete its shell makes it react vigorously. While fluorine bonds easily with other atoms, the resulting carbon–fluorine bonds are often very stable.

Fluorine shows different physical states depending on the temperature. At room temperature, it is a pale yellow gas with a pungent, biting odor. If you cool it, it condenses into a bright yellow liquid at -188.1°C.

It can also exist in two different solid forms: alpha-fluorine and beta-fluorine. Beta-fluorine is a transparent, soft solid with a disordered cubic structure. If you cool it even further, it undergoes a phase transition. It becomes alpha-fluorine, which is a hard, opaque solid with a monoclinic structure. This transition from beta to alpha is exothermic, meaning it releases heat, and can be quite violent.

Human history with fluorine began with the study of minerals. In 1529, Georgius Agricola described a mineral called fluorite. He used the Latin verb *fluorere*, meaning "to flow," to name it. This was because the mineral helped metal ores flow more easily during smelting. For centuries, people used fluorite to lower the melting points of metals. In 1810, the physicist André-Marie Ampère proposed the name "fluorine." However, isolating the pure element was incredibly difficult and dangerous. Many early researchers died or were injured while trying to separate it from its compounds. It was not until 1886 that Henri Moissan successfully isolated elemental fluorine.

He used a process called low-temperature electrolysis, which is still used today.

Fluorine is relatively rare in the universe. It ranks 24th in cosmic abundance, with a value of about 400 parts per billion. This low number is because fluorine is often bypassed during stellar nucleosynthesis. However, it is more common on Earth, where it ranks 13th in crustal abundance. Most fluorine on our planet is found in minerals like fluorite, fluorapatite, and cryolite.

China and Mexico are among the major global suppliers of these minerals. During World War II, the Manhattan Project began the industrial production of fluorine gas. This was necessary for the uranium enrichment process.

Today, the economic impact of fluorine is massive. Global sales of fluorochemicals exceed US$15 billion every year. Most commercial work uses fluorine compounds rather than the pure gas because refining the gas is expensive. About half of all mined fluorite is used in steelmaking. The rest is used to create hydrogen fluoride, which leads to various organic fluorides. Fluorine is also essential for making aluminium through the use of cryolite.

Some specialized gases, like sulfur hexafluoride (SF6), are used in electrical equipment.

You encounter fluorine in many unexpected places. In medicine, drugs like atorvastatin and fluoxetine contain carbon–fluorine bonds.

In your home, Teflon is a common organofluorine compound used for non-stick surfaces. Fluoride ions are also added to toothpaste and drinking water to prevent dental cavities.

However, there are environmental concerns to consider. Many fluorocarbon gases are potent greenhouse gases. SF6, for example, has a global warming potential up to 23,500 times higher than carbon dioxide. Because the carbon–fluorine bond is so strong, these compounds can persist in the environment for a long time.