FLUORINE
Fluorine is a chemical element with symbol F and atomic number 9. It is the lightest halogen and exists as a highly toxic pale yellow diatomic gas at standard conditions. As the most electronegative element, it is extremely reactive, as it reacts with almost all other elements, except for helium and neon.
Among the elements, fluorine ranks 24th in universal abundance and 13th in terrestrial abundance. Fluorite, the primary mineral source of fluorine which gave the element its name, was first described in 1529; as it was added to metal ores to lower their melting points for smelting, the Latin verb fluo meaning "flow" gave the mineral its name.
Reactions of Fluorine
Because of its reactivity, elemental fluorine is never found in nature and no other chemical element can displace fluorine from its compounds. Fluorine bonds with almost any element, both metals and nonmetals, because it is a very strong oxidizing agent. It is very unstable and reactive since it is so close to its ideal electron configuration. It forms covalent bonds with nonmetals, and since it is the most electronegative element, is always going to be the element that is reduced. It can also form a diatomic element with itself (F2F2), or covalent bonds where it oxidizes other halogens (ClFClF, ClF3ClF3, ClF5ClF5). It will react explosively with many elements and compounds such as Hydrogen and water. Elemental Fluorine is slightly basic, which means that when it reacts with water it forms OH−.
3F2 + 2H2O→O2 + 4HF
3F2 + 2H2O→O2 + 4HF
When combined with Hydrogen, Fluorine forms Hydrofluoric acid (HFHF), which is a weak acid. This acid is very dangerous and when dissociated can cause severe damage to the body because while it may not be painful initially, it passes through tissues quickly and can cause deep burns that interfere with nerve function.
HF + H2O→H3O+ + F−
HF + H2O→H3O+ + F−
There are also some organic compounds made of Fluorine, ranging from nontoxic to highly toxic. Fluorine forms covalent bonds with Carbon, which sometimes form into stable aromatic rings. When Carbon reacts with Fluorine the reaction is complex and forms a mixture of CF4CF4, C2F6C2F6, an C5F12C5F12.
C(s) + F2(g)→CF4(g) + C2F6 + C5F12
C(s) + F2(g)→CF4(g) + C2F6 + C5F12
Fluorine reacts with Oxygen to form OF2OF2 because Fluorine is more electronegative than Oxygen. The reaction goes:
2F2 + O2→2OF2
2F2 + O2→2OF2
Fluorine is so electronegative that sometimes it will even form molecules with noble gases like Xenon, such as the the molecule Xenon Difluoride, XeF2XeF2.
Xe + F2→XeF2
Xe + F2→XeF2
Fluorine also forms strong ionic compounds with metals. Some common ionic reactions of Fluorine are:
F2 + 2NaOH→O2 + 2NaF + H2
F2 + 2NaOH→O2 + 2NaF + H2
4F2 + HCl + H2O→3HF + OF2 + ClF3
4F2 + HCl + H2O→3HF + OF2 + ClF3
F2 + 2HNO3→2NO3F + H2
ISOTOPES
Only one isotope of fluorine occurs naturally in abundance, the stable isotope 19F. It has a high magnetogyric ratio and exceptional sensitivity to magnetic fields; because it is also the only stable isotope, it is used in magnetic resonance imaging.Seventeen radioisotopeswith mass numbers from 14 to 31 have been synthesized, of which 18F is the most stable with a half-life of 109.77 minutes. Other radioisotopes have half-lives less than 70 seconds; most decay in less than half a second. The isotopes 17F and 18F undergo β+ decayand electron capture, lighter isotopes decay by proton emission, and those heavier than 19F undergo β− decay (the heaviest ones with delayed neutron emission). Two metastable isomers of fluorine are known, 18mF, with a half-life of 162(7) nanoseconds, and 26mF, with a half-life of 2.2(1) milliseconds.
INDUSTRIAL APPLICATIONS
Fluorite mining, which supplies most global fluorine, peaked in 1989 when 5.6 million metric tons of ore were extracted. Chlorofluorocarbon restrictions lowered this to 3.6 million tons in 1994; production has since been increasing. Around 4.5 million tons of ore and revenue of US$550 million were generated in 2003; later reports estimated 2011 global fluorochemical sales at $15 billion and predicted 2016–18 production figures of 3.5 to 5.9 million tons, and revenue of at least $20 billion. Froth flotation separates mined fluorite into two main metallurgical grades of equal proportion: 60–85% pure metspar is almost all used in iron smelting whereas 97%+ pure acidspar is mainly converted to the key industrial intermediate hydrogen fluoride.
Among the elements, fluorine ranks 24th in universal abundance and 13th in terrestrial abundance. Fluorite, the primary mineral source of fluorine which gave the element its name, was first described in 1529; as it was added to metal ores to lower their melting points for smelting, the Latin verb fluo meaning "flow" gave the mineral its name.
- Allotropes: alpha, beta
- Appearance: GAS(very pale yellow), LIQUID(bright yellow), SOLID(alpha is opaque & beta is transparent)
- Standard Atomic Weight: 18.998
- Atomic Number(Z): 9
- Group: group 17 (halogens)
- Period: period 2
- Block: p block
- Element Category: reactive non-metal
- PHASE AT STP : gas
- Melting Point : 53.48 K [-219.67 ° C]
- Boiling Point : 85.03 K [-188.11 ° C]
- Density(at STP) : 1.696 g/L
- Heat of Vaporization : 6.51 kj/mol
- Covalent Radius : 64pm
- Crystal Structure : cubic crystal structure for fluorine
- Discovery : Andre-Marie Ampere(1810)
Reactions of Fluorine
Because of its reactivity, elemental fluorine is never found in nature and no other chemical element can displace fluorine from its compounds. Fluorine bonds with almost any element, both metals and nonmetals, because it is a very strong oxidizing agent. It is very unstable and reactive since it is so close to its ideal electron configuration. It forms covalent bonds with nonmetals, and since it is the most electronegative element, is always going to be the element that is reduced. It can also form a diatomic element with itself (F2F2), or covalent bonds where it oxidizes other halogens (ClFClF, ClF3ClF3, ClF5ClF5). It will react explosively with many elements and compounds such as Hydrogen and water. Elemental Fluorine is slightly basic, which means that when it reacts with water it forms OH−.
3F2 + 2H2O→O2 + 4HF
3F2 + 2H2O→O2 + 4HF
When combined with Hydrogen, Fluorine forms Hydrofluoric acid (HFHF), which is a weak acid. This acid is very dangerous and when dissociated can cause severe damage to the body because while it may not be painful initially, it passes through tissues quickly and can cause deep burns that interfere with nerve function.
HF + H2O→H3O+ + F−
HF + H2O→H3O+ + F−
There are also some organic compounds made of Fluorine, ranging from nontoxic to highly toxic. Fluorine forms covalent bonds with Carbon, which sometimes form into stable aromatic rings. When Carbon reacts with Fluorine the reaction is complex and forms a mixture of CF4CF4, C2F6C2F6, an C5F12C5F12.
C(s) + F2(g)→CF4(g) + C2F6 + C5F12
C(s) + F2(g)→CF4(g) + C2F6 + C5F12
Fluorine reacts with Oxygen to form OF2OF2 because Fluorine is more electronegative than Oxygen. The reaction goes:
2F2 + O2→2OF2
2F2 + O2→2OF2
Fluorine is so electronegative that sometimes it will even form molecules with noble gases like Xenon, such as the the molecule Xenon Difluoride, XeF2XeF2.
Xe + F2→XeF2
Xe + F2→XeF2
Fluorine also forms strong ionic compounds with metals. Some common ionic reactions of Fluorine are:
F2 + 2NaOH→O2 + 2NaF + H2
F2 + 2NaOH→O2 + 2NaF + H2
4F2 + HCl + H2O→3HF + OF2 + ClF3
4F2 + HCl + H2O→3HF + OF2 + ClF3
F2 + 2HNO3→2NO3F + H2
ISOTOPES
Only one isotope of fluorine occurs naturally in abundance, the stable isotope 19F. It has a high magnetogyric ratio and exceptional sensitivity to magnetic fields; because it is also the only stable isotope, it is used in magnetic resonance imaging.Seventeen radioisotopeswith mass numbers from 14 to 31 have been synthesized, of which 18F is the most stable with a half-life of 109.77 minutes. Other radioisotopes have half-lives less than 70 seconds; most decay in less than half a second. The isotopes 17F and 18F undergo β+ decayand electron capture, lighter isotopes decay by proton emission, and those heavier than 19F undergo β− decay (the heaviest ones with delayed neutron emission). Two metastable isomers of fluorine are known, 18mF, with a half-life of 162(7) nanoseconds, and 26mF, with a half-life of 2.2(1) milliseconds.
INDUSTRIAL APPLICATIONS
Fluorite mining, which supplies most global fluorine, peaked in 1989 when 5.6 million metric tons of ore were extracted. Chlorofluorocarbon restrictions lowered this to 3.6 million tons in 1994; production has since been increasing. Around 4.5 million tons of ore and revenue of US$550 million were generated in 2003; later reports estimated 2011 global fluorochemical sales at $15 billion and predicted 2016–18 production figures of 3.5 to 5.9 million tons, and revenue of at least $20 billion. Froth flotation separates mined fluorite into two main metallurgical grades of equal proportion: 60–85% pure metspar is almost all used in iron smelting whereas 97%+ pure acidspar is mainly converted to the key industrial intermediate hydrogen fluoride.
