FLUORINE SYMBOL:F PERIOD:2 GROUP:17(VIIA) ATOMICNO:9
FLUORINE SYMBOL:F PERIOD:2 GROUP:17(VIIA) ATOMICNO:9
ATOMICMASS:18.99840amu VALENCE:1 OXIDATIONSTATE:–2 NATURALSTATE:Gas ORIGINOFNAME: FromtheLatinandFrenchwordsfor“flow,”fluere. ISOTOPES:Thereareatotalof16isotopesoffluorine.Onlyone,F-19,isstable.Itmakes
up100%ofthefluorinefoundonEarth.Alltheothersareradioactivewithhalf-lives
246 | The History and Use of Our Earth’s Chemical Elements
ELECTRONCONFIGURATION EnergyLevels/Shells/Electrons Orbitals/Electrons
s2,p5
Properties Fluorine does not occur in a free state in nature, and because fluorine is one of the most
reactive elements, no chemical can free it from any of its many compounds. The reason for this is that fluorine atoms are the smallest of the halogens, meaning the electron donated by a metal (or some nonmetals) are closer to fluorine’s nucleus and thus exert a great force between the fluorine nuclei and the elements giving up one electron. The positive nuclei of fluorine have a strong tendency to gain electrons to complete the outer shell, which makes it a strong oxidizer.
Because the fluorine atom has only nine electrons, which are close to the nucleus, the positive nucleus has a strong tendency to gain electrons to complete its outer shell. As a gas its density (specific gravity) is 1.695, and as a liquid, its density is 1.108. Its freezing point is
–219.61°C, and its boiling point is –188°C. Fluorine, as a diatomic gas molecule (F 2 ), is pale yellow in color. Fluorine is the most electronegative nonmetallic element known (wants to gain electrons) and is, therefore, the strongest oxidizing agent known.
Characteristics Fluorine reacts violently with hydrogen compounds, including water and ammonia. It also
reacts with metals, such as aluminum, zinc, and magnesium, sometimes bursting into flames, and with all organic compounds, in some cases resulting in such complex fluoride compounds as fluorocarbon molecules. It is an extremely active, gaseous element that combines spontane- ously and explosively with hydrogen, producing hydrogen fluoride acid (HF), which is used to etch glass. It reacts with most metals except helium, neon, and argon. It forms many dif- ferent types of “salts” when combining with a variety of metals. Fluorine, as a diatomic gas, is extremely poisonous and irritating to the skin and lungs, as are many fluoride compounds. Fluorine and its compounds are also corrosive.
AbundanceandSource Fluorine is the 13th most abundant element on the Earth. It makes up about 0.06% of the
Earth’s crust. Fluorine is widely distributed in many types of rocks and minerals, but never found in its pure form. Fluorine is as plentiful as nitrogen, chlorine, and copper, but less plentiful than aluminum or iron.
The most abundant fluorine mineral is fluorite—calcium fluoride (CaF 2 )—which is often found with other minerals, such as quartz, barite, calcite, sphalerite, and galena. It is mined in
247 Cumberland, England, and in Illinois in the United States. Other minerals from which fluo-
Guide to the Elements |
rine is recovered are fluorapatite, cryolite, and fluorspar, which are found in many countries but mainly in Mexico and Africa.
Today fluorine is produced by the electrolysis of potassium fluoride (KF), hydrofluoric acid
(HF), and molten potassium acid fluoride (KHF 2 ).
History Fluorine was mentioned first in history in 1670 when instructions were written regarding its
use to etch glass, using green fluorspar (fluorite), which is calcium fluoride (CaF 2 ). In the early 1700s chemists tried to identify the material that etched glass. Although Carl Wilhelm Scheele first “discovered” fluorine in 1771, he was not given credit because the element was not yet isolated and correctly identified. In 1869 George Gore produced a small amount of fluorine through an electrolytic process. He was unaware that fluorine gas would react with the hydrogen produced at the other electrode and would become extremely explosive (his apparatus exploded). In 1886 a French chemist, Ferdinand Frederich Henri Moissan (1852–1907), used platinum electrodes to produce fluorine from the electrolysis of potassium fluoride (KF) and hydrofluoric acid (HF). He was able to contain each of the gasses separately, thus preventing an explosion. Moissan was credited with the discovery of fluorine, partly because of his unique way of produc- ing and identifying the element. He was awarded the 1906 Nobel Prize for Chemistry.
CommonUses Probably the most common use of fluorine is its addition to municipal water supplies to
help prevent tooth decay. Stannous (II) fluoride (SnF 2 ) is added to the water in proportions of about one part per million (1 ppm). In addition, many brands of toothpaste add stannous fluoride or other fluoride compounds to their product to help prevent tooth decay. Tooth enamel degenerates overtime. Fluorine promotes remineralization, essentially making a form of new enamel called “fluorapatite,” which is resistant to decay.
Another popular use for the element fluorine is the plastic called Teflon. This is a fluo- ropolymer consisting of long chainlike inert molecules of carbon linked chemically to fluorine. Teflon is useful as a coating for nonstick surfaces in cookware, ironing board covers, razor blades, and so forth.
Of great importance are the inert fluorocarbons, such as dichlorodifluoromethane (CF 2 Cl 2 ) and chlorofluorocarbon compounds (CFCs) and their usage as gas propellants in spray cans (e.g., hair spray, deodorants, and paint). They are also used as coolants in air conditioning and refrigeration (freon). The use of fluorinated carbon gases, known as fluorocarbons, in aerosol cans and refrigerants has been banned in the United States since 1978 because these gases dif- fuse into the upper atmosphere and react to destroy the ozone gases found in the ozone layer.
A reduced ozonosphere layer allows more ultraviolet radiation to filter to the Earth’s surface. Excessive strong ultraviolet radiation from the sun can be harmful to both plants and animals. The ozone layer filters out most of the harmful ultraviolet radiation
When hydrogen and fluorine gases meet, they explode spontaneously and form hydrogen fluoride (HF), which, when dissolved in water, becomes hydrofluoric acid that is strong enough to dissolve glass. It is used to etch glass and to produce “frosted” light bulbs.
The artificial radioactive fluorine isotope F-18 emits positrons (positive electrons) that, when injected into the body, interact with regular negative electrons, and they annihilate
248 | The History and Use of Our Earth’s Chemical Elements each other, producing X-ray-like radiation. This medical procedure is performed in Positron
Emission Topography (PET), in which the produced radiation generates a picture of the body part being examined. Since F-18 has a short half-life of about 110 minutes, there is little chance of radiation damage to the patient.
Fluorine compounds are also used to reduce the viscosity of molten metals and slag by- products so that they will flow more easily. In addition, fluorine is a component of therapeutic chemotherapy drugs used to treat a number of different types of cancer.
ExamplesofCompounds Hydrogen fluoride (hydrofluoric acid) is commercially prepared by distilling a mixture of
calcium fluoride (feldspar) with concentrated sulfuric acid, as follows: CaF 2 +H 2 SO 4 → 2HF + CaSO 4.
Fluorine nitrate (FNO 3 ) is a strong oxidizing gas or liquid. In the liquid state it explodes by shock or friction. It is used as an oxidizer for rocket propellant fuels. Fluoroacetic acid (CH 2 FCOOH) is very poisonous. It is used to kill rats and mice. Chlorofluorocarbons (CFCs) come in many forms, including those used as propellants for spray cans and for refrigeration (freon). They were banned as being potentially harmful to the ozone layer of the atmosphere. In 1987 an international agreement was signed by about 90 nations to reduce the use of CFCs by 50% by the year 2000. This did not seem adequate, so in 1990 a new treaty called for the elimination of the use of all CFCs by industrial nations. Some third world countries (e.g., China, India, Russia, and Mexico) still make and sell CFCs, some of which are smuggled into the United States.
Sodium fluoride (NaF), in the concentration of one ppm, is added to municipal drinking water to help reduce tooth decay. It is also used as an insecticide, fungicide, and rodenticide, as well as in the manufacture of adhesives, disinfectants, and dental products.
Hazards Many of the fluorine compounds, such as CFCs, are inert and nontoxic to humans. But
many other types of compounds, particularly the salts and acids of fluorine, are very toxic when either inhaled or ingested. They are also strong irritants to the skin.
There is also danger of fire and explosion when fluorine combines with several elements and organic compounds.
Poisonous fluoride salts are not toxic to the human body at the very low concentration levels used in drinking water and toothpaste to prevent dental decay.
CHLORINE SYMBOL:Cl PERIOD:3 GROUP:17(VIIA) ATOMICNO:17
ATOMICMASS:35.453amu VALENCE:1,3,4,5,and7 OXIDATIONSTATE:–1 NAT-
URALSTATE:Gas ORIGINOFNAME: FromtheGreekwordkhlôros,meaning“greenishyellow.” ISOTOPES:Thereareatotalof25isotopesofchlorine.Ofthese,onlytwoarestableand
contributetothenaturalabundanceonEarthasfollows:Cl-35=75.77%andCl-37= 24.23%.Alltheother23isotopesareproducedartificially,areradioactive,andhavehalf- livesrangingfrom20nanosecondsto3.01×10 +5 years.
Guide to the Elements |
ELECTRONCONFIGURATION EnergyLevels/Shells/Electrons Orbitals/Electrons
s2,p6
3-M=7
s2,p5
Properties As a nonmetal, chlorine exists as a greenish-yellow gas that is corrosive and toxic at room
temperatures. As a halogen, chlorine is not found in the elemental (atomic) state but forms diatomic gas molecules (Cl ). As a very active negative ion with the oxidation state of –1, 2 chlorine forms bonds with most metals found in groups I and II.
Chlorine is noncombustible but will support combustion. It is extremely electronegative and a strong oxidizing agent. It is not as strong as fluorine, which is just above it in group 17, but is stronger than the other halogens.
As a gas, its specific gravity (density) is 3.214g/l or 0.003214g/cm 3 . As a liquid, it is a clear amber color with a density of 1.56g/cm 3 . Its melting point is –101.5°C, and its boiling point is –34.04°C.