Fluorine phase at room temperature

Fluorine phase at room temperature DEFAULT

Atomic Number: 9

Atomic Weight:

Melting Point: K (°C or °F)

Boiling Point: K (°C or °F)

Density: grams per cubic centimeter

Phase at Room Temperature: Gas

Element Classification: Non-metal

Period Number: 2

Group Number: 17

Group Name: Halogen

What's in a name? From the Latin and French words for flow, fluere.

Say what? Fluorine is pronounced as FLU-eh-reen or as FLU-eh-rin.

History and Uses:

Fluorine is the most reactive of all elements and no chemical substance is capable of freeing fluorine from any of its compounds. For this reason, fluorine does not occur free in nature and was extremely difficult for scientists to isolate. The first recorded use of a fluorine compound dates to around to a set of instructions for etching glass that called for Bohemian emerald (CaF2). Chemists attempted to identify the material that was capable of etching glass and George Gore was able to produce a small amount of fluorine through an electrolytic process in Unknown to Gore, fluorine gas explosively combines with hydrogen gas. That is exactly what happened in Gore's experiment when the fluorine gas that formed on one electrode combined with the hydrogen gas that formed on the other electrode. Ferdinand Frederic Henri Moissan, a French chemist, was the first to successfully isolate fluorine in He did this through the electrolysis of potassium fluoride (KF) and hydrofluoric acid (HF). He also completely isolated the fluorine gas from the hydrogen gas and he built his electrolysis device completely from platinum. His work was so impressive that he was awarded the Nobel Prize for chemistry in Today, fluorine is still produced through the electrolysis of potassium fluoride and hydrofluoric acid as well as through the electrolysis of molten potassium acid fluoride (KHF2).

Fluorine is added to city water supplies in the proportion of about one part per million to help prevent tooth decay. Sodium fluoride (NaF), stannous(II) fluoride (SnF2) and sodium monofluorophosphate (Na2PO3F) are all fluorine compounds added to toothpaste, also to help prevent tooth decay. Hydrofluoric acid (HF) is used to etch glass, including most of the glass used in light bulbs. Uranium hexafluoride (UF6) is used to separate isotopes of uranium. Crystals of calcium fluoride (CaF2), also known as fluorite and fluorspar, are used to make lenses to focus infrared light. Fluorine joins with carbon to form a class of compounds known as fluorocarbons. Some of these compounds, such as dichlorodifluoromethane (CF2Cl2), were widely used in air conditioning and refrigeration systems and in aerosol spray cans, but have been phased out due to the damage they were causing to the earth's ozone layer.

Estimated Crustal Abundance: ×102 milligrams per kilogram

Estimated Oceanic Abundance: milligrams per liter

Number of Stable Isotopes: 1 (View all isotope data)

Ionization Energy: eV

Oxidation States: -1

Sours: https://education.jlab.org/itselemental/elehtml

Phases of fluorine

Fluorine forms diatomic molecules (F
2) that are gaseous at room temperature with a density about times that of air.[note 1] Though sometimes cited as yellow-green, pure fluorine gas is actually a very pale yellow. The color can only be observed in concentrated fluorine gas when looking down the axis of long tubes, as it appears transparent when observed from the side in normal tubes or if allowed to escape into the atmosphere.[3] The element has a "pungent" characteristic odor that is noticeable in concentrations as low as 20&#;ppb.[citation needed]

observation of the color of fluorine gas by Henri Moissan, who first isolated the element, using end-view on 5 meter long tubes. Air (1) is on the left, fluorine (2) is in the middle, chlorine (3) is on the right.

Fluorine condenses to a bright yellow liquid at −&#;°C (−&#;°F), which is near the condensation temperatures of oxygen and nitrogen.

The solid state of fluorine relies on Van der Waals forces to hold molecules together,[citation needed] which, because of the small size of the fluorine molecules, are relatively weak. Consequently, the solid state of fluorine is more similar to that of oxygen or the noble gases than to those of the heavier halogens.[citation needed]

Animation showing the crystal structure of beta-fluorine.

Fluorine solidifies at −&#;°C (−&#;°F) into a cubic structure, called beta-fluorine. This phase is transparent and soft, with significant disorder of the molecules; its density is &#;g/cm3. At −&#;°C (−&#;°F) fluorine undergoes a solid–solid phase transition into a monoclinic structure called alpha-fluorine. This phase is opaque and hard, with close-packed layers of molecules, and is denser at &#;g/cm3.[5] The solid state phase change requires more energy than the melting point transition and can be violent, shattering samples and blowing out sample holder windows.[6][7]

Solid fluorine received significant study in the s and 30s, but relatively less until the s. The crystal structure of alpha-fluorine given, which still has some uncertainty, dates to a paper by Linus Pauling.

angled lines showing linear pressure temp relations of the lower phase boundariesA parallelogram-shaped outline with space-filling diatomic molecules (joined circles) arranged in two layers
Low-temperature fluorine phases Alpha-fluorine crystal structure

Notes[edit]

Citations[edit]

Indexed references[edit]

Further reading[edit]

  • Jordan, T. H.; Streib, W. E.; Lipscomb, W. N. (). "Single-Crystal X-Ray Diffraction Study of β-Fluorine". The Journal of Chemical Physics. 41 (3): BibcodeJChPhJ. doi/
  • Jordan, T. H.; Streib, W. D.; Smith, H. W.; Lipscomb, W. N. (). "Single-crystal studies of β-F2and of γ-O2". Acta Crystallographica. 17 (6): doi/SXX.
  • Meyer, L. (). "Crystal Structure of α-Fluorine". The Journal of Chemical Physics. 49 (4): – BibcodeJChPhM. doi/
  • Pauling, L.; Keaveny, I.; Robinson, A. B. (). "The Crystal Structure of α-Fluorine". Journal of Solid State Chemistry. 2 (2): – BibcodeJSSChP. doi/(70)
  • Etters, R. D.; Kirin, D. (). "High-pressure behavior of solid molecular fluorine at low temperatures". The Journal of Physical Chemistry. 90 (19): doi/ja
  • Kobashi, K.; Klein, M. L. (). "Lattice vibrations of solid α-F2". Molecular Physics. 41 (3): BibcodeMolPhK. doi/
  • English, C. A.; Venables, J. A. (). "The Structure of the Diatomic Molecular Solids". Proceedings of the Royal Society A: Mathematical, Physical and Engineering Sciences. (): BibcodeRSPSAE. doi/rspa
  • http://www.osti.gov/bridge/servlets/purl/BbwUC/pdf (phase diagrams of the elements)
  • http://jcp.aip.org/resource/1/jcpsa6/v47/i2/p_s1?isAuthorized=no (sample holder blowout)
  • NASA ADS: Solid Fluorine and Solid Chlorine: Crystal Structures and Intermolecular Forces by S. C. Nyburg
Sours: https://en.wikipedia.org/wiki/Phases_of_fluorine
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In the periodic table above, black squares indicate elements which are solids at room temperature (about 22ºC)*, those in blue squares are liquids at room temperature, and those in red squares are gases at room temperature.

Most of the metals are solids under "ordinary" conditions (i.e., 25ºC, 1 atmosphere of pressure, etc.), with the exception of mercury (Hg, element 80), which solidifies at ºC, and is a freely-flowing liquid at room temperature.  Gallium (Ga, element 31) melts at 30ºC, slightly above room temperature, but is often indicated as a liquid on periodic tables, since the solid metal literally melts when held in the hand (since body temperature is about 37ºC).  (Since cesium melts at 28ºC, and francium at 27ºC, they are also indicated in blue on some tables, but anyone who holds cesium in their hands won't be holding much of anything afterwards!  See the page on alkali metals for more on cesium's high reactivity.)

Several of the nonmetals are gases in their elemental form.  Elemental hydrogen (H, element 1), nitrogen (N, element 7), oxygen (O, element 8), fluorine (F, element 9), and chlorine (Cl, element 17) are all gases at room temperature, and are found as diatomic molecules (H2, N2, O2, F2, Cl2).  Bromine (Br, element 35), also found as a diatomic molecule (Br2), is a liquid at room temperature, solidifying at ºC.  The noble gases of Group 8A (He, Ne, Ar, Kr, Xe, and Rn) are all gases at room temperature (as the name of the group implies); since they are all unreactive, monatomic elements, their boiling points are extremely low.

Below is a table of the melting points, boiling points, and densities of the elements:

 

Atomic
Number

Symbol

Name

Melting
Point
(ºC)

Boiling
Point
(ºC)

Density
(g/cm3)
(at K)

1

H

Hydrogen

(gas, K)

2

He

Helium

(under pressure) (gas, K)

3

Li

Lithium

4

Be

Beryllium

5

B

Boron

6

C

Carbon


(sublimes)
(graphite)
(diamond)

7

N

Nitrogen

(gas, K)

8

O

Oxygen

(gas, K)

9

F

Fluorine

(gas, K)

10

Ne

Neon

(gas, K)

11

Na

Sodium

12

Mg

Magnesium

13

Al

Aluminum

14

Si

Silicon

15

P

Phosphorus

(white)
(red, under pressure)
(white) (white)

16

S

Sulfur

(&#;)
(b)
(g)
(&#;)
(b)

17

Cl

Chlorine

(gas, K)

18

Ar

Argon

(gas, K)

19

K

Potassium

20

Ca

Calcium

21

Sc

Scandium

22

Ti

Titanium

23

V

Vanadium

( K)

24

Cr

Chromium

25

Mn

Manganese

26

Fe

Iron

27

Co

Cobalt

28

Ni

Nickel

( K)

29

Cu

Copper

30

Zn

Zinc

31

Ga

Gallium

32

Ge

Germanium

33

As

Arsenic

(under pressure) (sublimes) (&#;)

34

Se

Selenium

35

Br

Bromine

36

Kr

Krypton

(gas, K)

37

Rb

Rubidium

38

Sr

Strontium

39

Y

Yttrium

40

Zr

Zirconium

41

Nb

Niobium

42

Mo

Molybdenum

43

Tc

Technetium

(est.)

44

Ru

Ruthenium

45

Rh

Rhodium

46

Pd

Palladium

47

Ag

Silver

48

Cd

Cadmium

49

In

Indium

( K)

50

Sn

Tin

(&#;)
(b)

51

Sb

Antimony

52

Te

Tellurium

53

I

Iodine

54

Xe

Xenon

(gas, K)

55

Cs

Cesium

56

Ba

Barium

57

La

Lanthanum

( K)

58

Ce

Cerium

(a, K)

59

Pr

Praseodymium

60

Nd

Neodymium

61

Pm

Promethium

ca. ( K)

62

Sm

Samarium

63

Eu

Europium

64

Gd

Gadolinium

( K)

65

Tb

Terbium

66

Dy

Dysprosium

67

Ho

Holmium

( K)

68

Er

Erbium

( K)

69

Tm

Thulium

70

Yb

Ytterbium

71

Lu

Lutetium

( K)

72

Hf

Hafnium

73

Ta

Tantalum

74

W

Tungsten

75

Re

Rhenium

76

Os

Osmium

77

Ir

Iridium

( K)

78

Pt

Platinum

79

Au

Gold

80

Hg

Mercury

81

Tl

Thallium

82

Pb

Lead

83

Bi

Bismuth

84

Po

Polonium

85

At

Astatine

n/a

86

Rn

Radon

(gas, K)

87

Fr

Francium

27n/a

88

Ra

Radium

89

Ac

Actinium

90

Th

Thorium

91

Pa

Protactinium

92

U

Uranium

93

Np

Neptunium

94

Pu

Plutonium

(&#;, K)

95

Am

Americium

96

Cm

Curium

n/a

97

Bk

Berkelium

n/a

98

Cf

Californium

n/an/a

99

Es

Einsteinium

n/an/a

Fm

Fermium

n/an/an/a

Md

Mendelevium

n/an/an/a

No

Nobelium

n/an/an/a

Lr

Lawrencium

n/an/an/a

Rf

Rutherfordium

(est.) (est.)23 (est.)

Db

Dubnium

n/an/a29

Sg

Seaborgium

n/an/a35 (est.)

Bh

Bohrium

n/an/a37 (est.)

Hs

Hassium

n/an/a41

Mt

Meitnerium

n/an/an/a

Uun

Ununnilium

n/an/an/a

Uuu

Unununium

n/an/an/a

Uub

Ununbiium

n/an/an/a

——

———

———

———

Uuq

Ununquadium

n/an/an/a

Data taken from John Emsley, The Elements, 3rd edition.  Oxford:  Clarendon Press,

 

 

 

 

 

 

* It doesn't matter what temperature a room is; it's always room temperature.

Stephen Wright

Sours: https://www.angelo.edu/faculty/kboudrea/periodic/physical_states.htm
Fluorine

Transcript :

Chemistry in its element: fluorine

(Promo)

You're listening to Chemistry in its element brought to you by Chemistry World, the magazine of the Royal Society of Chemistry.

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Chris Smith

This week, a strong acid it's not, but deadly it definitely is.

Kira J. Weissman

The year old technician spilled only a few hundred milliliters or so in his lap during a routine palaeontology experiment. He took the normal precaution in such situations, quickly dowsing himself with water from a laboratory hose, and even plunged into a nearby swimming pool while the paramedics were en route. But a week later, doctors removed a leg, and a week after that, he was dead. The culprit: hydrofluoric acid (colloquially known as HF), and the unfortunate man was not its first victim.

Unlike its close relatives, hydrochloric and hydrobromic acid, HF is a weak acid. This, coupled with its small molecular size, allows it to penetrate the skin and migrate rapidly towards the deeper tissue layers. Once past the epidermis, HF starts to dissociate, unleashing the highly-reactive fluoride ion. Free fluoride binds tightly to both calcium and magnesium, forming insoluble salts which precipitate into the surrounding tissues. Robbed of their co-factors, critical metabolic enzymes can no longer function, cells begin to die, tissues to liquefy and bone to corrode away. And if calcium loss is rapid enough, muscles such as the heart stop working. Burns with concentrated HF involving as little as % of the body surface area - the size of the sole of the foot, for example - have been fatal.

HF has a long history of destructive behaviour, claiming the lives of several chemists in the s, including the Belgian Paulin Louyet, and the Frenchman Jérôme Nicklès. These brave scientists were battling to be the first to isolate elemental fluorine (F2) from its various compounds, using electrolysis. However, it was Nicklès' countrymen, Henri Moissan, who succeeded in To achieve this feat, Moissan not only had to contend with HF - the preferred electrolyte in such experiments - but fluorine itself, a violently reactive gas. His key innovation was to construct an apparatus out of platinum, one of the few metals capable of resisting attack, while cooling the electrolytic solution down to °C to limit corrosion. Moissan's feat earned him the Nobel Prize in chemistry, but the celebration was short-lived. Another victim of fluorine's toxic effects, he died only two months later. Yet Moissan's method lived on, and is used today to produce multi-ton quantities of fluorine from its ore fluorspar.

Ironically, while elemental fluorine is decidedly bad for your health, fluorine atoms turns up in some 20% of all pharmaceuticals. The top-selling anti-depressant Prozac, the cholesterol-lowering drug Lipitor, and the antibacterial Cipro, all have fluorine to thank for their success. How is this possible? Because the flip side of fluorine's extreme reactivity is the strength of the bonds it forms with other atoms, notably including carbon. This property makes organofluorine compounds some of the most stable and inert substances known to man. Fluorine's special status also stems from the 'fluorine factor', the ability of this little atom to fine-tune the chemical properties of an entire molecule. For example, replacing hydrogen with fluorine can protect drugs from degradation by metabolic enzymes, extending their active lifetimes inside the body. Or the introduced fluorine can alter a molecule's shape so that it binds better to its target protein. Such precise chemical tinkering can now be carried out in pharmaceutical labs using an array of safe, commercially-available fluorinating agents, or the tricky transformations can simply be out-sourced to someone else.

Most of us also have fluorine to thank for our beaming smiles. The cavity-fighting agents in toothpaste are inorganic fluorides such as sodium fluoride and sodium monofluorophosphate. Fluoride not only decreases the amount of enamel-dissolving acid produced by plaque bacteria, but aids in the tooth rebuilding process, insinuating itself into the enamel to form an even harder surface which resists future attack. And the list of medical applications doesn't stop there. Being put to sleep is a little bit less worrisome thanks to fluorinated anaesthetics such as isoflurane and desflurane, which replaced flammable and explosive alternatives such as diethyl ether and chloroform. Fluorocarbons are also one of the leading candidates in development as artificial blood, as oxygen is more soluble in these materials than most other solvents. And radioactive fluorine (18F rather than the naturally-occurring 19F) is a key ingredient in positron emission tomography (or PET), a whole-body imaging technique that allows cancerous tumours to be discovered before they spread.

Fluorochemicals are also a mainstay of industry. One of the most famous is the polymer polytetrafluoroethylene, better known as Teflon, which holds the title of world's most slippery solid. Highly thermostable and water proof, it's used as a coating for pots and pans, in baking sprays, and to repel stains on furniture and carpets. Heating and stretching transforms Teflon into Gore-tex, the porous membrane of sportswear fame. Gore-tex's pores are small enough to keep water droplets out, while allowing water vapour (that is, sweat) to escape. So you can run on a rainy day, and still stay dry. Fluorine plays another important role in keeping you cool, as air-conditioning and household refrigeration units run on energy-efficient fluorocarbon fluids. And fluorine's uses are not limited to earth. When astronauts jet off into space they put their trust in fluoroelastomers, a type of fluorinated rubber. Fashioned into O-rings and other sealing devices, these materials ensure that aircraft remain leak-free even under extreme conditions of heat and cold. And when accidents do happen, space travellers can rely on fluorocarbon-based fire extinguishers to put the flames out.

Fluorine has long been known as the 'tiger of chemistry'. And while the element certainly retains its wild side, we can reasonably claim to have tamed it. As only a handful of naturally-occurring organofluorine compounds have ever been discovered, some might argue that we now make better use of fluorine than even Nature herself.

Chris Smith

So Teflon is acknowledged as the world's most slippery thing and I bet there are one or two politicians knocking around who are thanking fluorine for that. Thank you also to Kira Weismann from Zaarland University in Germany. Next week.ouch

Steve Mylon

I cannot imagine that this is all someone would be saying if they were unfortunate enough to be stricken with the disease of the same name. The ouch-ouch disease.

The disease results from excessive cadmium poisoning and was first reported in a small town about miles north west of Tokyo. Rice grown in cadmium contaminated soils had more than 10 times the cadmium content than normal rice. The ouch-ouch-ness of this disease resulted from weak and brittle bones subject to collapse due to high porosity.

Chris Smith

And you can find out about the ouch-ouch factor with Steve Mylon when he uncovers the story of cadmium on next week's Chemistry in Its Element. I'm Chris Smith, thank you for listening and goodbye.

(Promo)

Chemistry in its element is brought to you by the Royal Society of Chemistry and produced by thenakedscientists.com. There's more information and other episodes of Chemistry in its element on our website at chemistryworld.org/elements.

(End promo)

Sours: https://www.rsc.org/periodic-table/element/9/fluorine

Phase temperature room fluorine at

Predicting properties

The halogens show trends in physical properties as you go down the group.

Melting point and boiling point

The halogens have low melting points and low boiling points. This is a typical property of non-metals. Fluorine has the lowest melting and boiling points. The melting and boiling points then increase as you go down the group.

State at room temperature

Room temperature is usually taken as being 25°C. At this temperature, fluorine and chlorine are gases, bromine is a liquid, and iodine and astatine are solids. There is therefore a trend in state from gas to liquid to solid as you go down the group.

Colour

The halogens become darker as you go down the group. Fluorine is very pale yellow, chlorine is yellow-green, and bromine is red-brown. Iodine crystals are shiny purple - but easily turn into a dark purple vapour when they are warmed up.

Predictions

When we can see a trend in the properties of some of the elements in a group, it is possible to predict the properties of other elements in that group. Astatine is below iodine in Group 7. The colour of these elements gets darker as you go down the group. Iodine is purple, and astatine is black.

Sours: https://www.bbc.co.uk/bitesize/guides/ztq6cwx/revision/2
Why is Oxygen (O2) a gas at room temperature while Water (H2O) is a liquid?

Facts About Fluorine

The most reactive element on the Periodic Table, fluorine has a violent history in the quest for its discovery. Despite the difficult and sometimes explosive properties of fluorine, it is a vital element for humans and animals, which is why it is commonly found in drinking water and toothpaste.

Just the facts

  • Atomic number (number of protons in the nucleus): 9
  • Atomic symbol (on the Periodic Table of Elements): F
  • Atomic weight (average mass of the atom):
  • Density: grams per cubic centimeter
  • Phase at room temperature: Gas
  • Melting point: minus degrees Fahrenheit (minus degrees Celsius)
  • Boiling point: minus degrees F (minus degrees C)
  • Number of isotopes (atoms of the same element with a different number of neutrons): 18
  • Most common isotopes: F ( percent natural abundance)

History

Early chemists tried for years to isolate the element from various fluorides. It wasn't until that German chemist Karl O. Christie successfully synthesized fluorine, and reported his results in the journal Inorganic Chemistry.  Fluorine does not occur free in nature; but in , researchers found small amounts of fluorine trapped in  antozonite, a type of radioactive fluorite.  

For centuries, the mineral fluorspar was used in metal refining. Known today as calcium fluoride (CaF2), it was used as a flux to separate pure metal from the unwanted minerals in ore, according to Chemicool. The "fluor" comes from the Latin word "fluere," meaning "to flow," because that's what fluorspar allowed metals to do. The mineral was also called Bohemian emerald and was used in glass etching, according to the Jefferson Laboratory.

Many scientists over the decades attempted to experiment with fluorspar to better learn its properties, as well as its composition. In their experiments, chemists often produced fluoric acid (today known as hydrofluoric acid, HF), an incredibly reactive and dangerous acid. Even small splashes of this acid on skin can be fatal, according to Chemicool. Several scientists were injured, blinded or killed in some of the experiments.

In the early 19th century, scientists Andre-Marie Ampere, in France, and Humphry Davy, in England, corresponded about the possibility of a new element within the acid. In , Davy announced the discovery of the new element and named it fluorine from Ampere's suggestion.

Henri Moissan, a French chemist, finally isolated fluorine in — after being poisoned several times in his pursuit. He was awarded the Nobel Prize in for the isolating fluorine by electrolysis of dry potassium hydrogen fluoride (KHF2) and dry hydrofluoric acid.

Uses of fluorine

For many years, fluorine salts, or fluorides, have been used in welding and for frosting glass, according to the Royal Society. For example, hydrofluoric acid is used to etch the glass of light bulbs.

Fluorine is a vital element in the nuclear energy industry, according to the Royal Society. It is used to make uranium hexafluoride, which is needed to separate uranium isotopes. Sulfur hexafluoride is a gas used to insulate high-power electricity transformers.

Chlorofluorocarbons (CFCs) were once used in aerosols, refrigerators, air conditioners, foam food packaging, and fire extinguishers. Those uses have been banned since because they contribute to ozone depletion, according to the National Institutes of Health. Prior to , CFCs were used in inhalers to control asthma but those types of inhalers were phased out in  

Fluorine is used in many fluorochemicals, including solvents and high-temperature plastics, such as Teflon (poly(tetrafluoroethene), PTFE). Teflon is well known for its non-stick properties and is used in frying pans. It is also used for cable insulation, for plumber's tape and as the basis of Gore-Tex® (used in waterproof shoes and clothing).

Fluorine is added to city water supplies in the proportion of about one part per million to help prevent tooth decay, according to the Jefferson Lab. Several fluoride compounds are added to toothpaste, also to help prevent tooth decay.

Health and environmental impacts of fluorine

Although all humans and animals are exposed to and need minute amounts of fluorine, the element in any large enough dose is extremely toxic and dangerous. According to Lenntech, fluorine can naturally be found in water, air, and both plant and animal-based foods in small amounts. Larger amounts of fluorine are found in a few food products such as tea and shellfish. 

While small amounts of fluorine are essential for maintaining the strength of our bones and teeth, too much can have the reverse effect of causing osteoporosis or tooth decay, as well has potentially harming the kidneys, nerves, and muscles.

In its gaseous form, fluorine is incredibly dangerous. Small amounts of fluorine gas can case eye and nose irritation while larger amounts can be fatal, according to Lenntech. Hydrofluoric acid, as another example, can also prove to be fatal when even a small splash on the skin occurs, according to Chemicool.

In the environment, fluorine, the 13th most abundant element in Earth's crust, typically settles within the soil and readily combines with soil, rock, coal and clay, according to Lenntech. Plants may absorb the fluorine from the soil, although high concentrations can lead to damage. Corn and apricots, for example, are among the plants that are most susceptible to damage and growth reduction when exposed to elevated levels of fluorine.

Who knew?

  • Because fluorine is the most chemically reactive element, it must be handled with extreme care as it can sometimes explode on contact with all other elements excluding oxygen, helium, neon and krypton, according to Chemicool.
  • Steel wool bursts into flames when exposed to fluorine, according to the Royal Society of Chemistry.
  • Fluorine is also the most electronegative element. Fluorine attracts electrons more readily than any other element.
  • On average, the amount of fluorine in the human body is three milligrams.
  • Fluorine is primarily mined in China, Mongolia, Russia, Mexico and South Africa, according to Minerals Education Coalition.
  • Fluorine is created in sun-like stars towards the end of their lifetime, according to a article published in the Astrophysical Journal Letters. The element is formed under the higher pressures and temperatures within the star when it expands to become a red giant. When the outer layers of the star are pushed away creating a planetary nebula, the fluorine travels along with the other gases into the interstellar medium eventually forming new stars and planets.
  • According to the Journal of Chemistry, approximately 25 percent of drugs and medications, including those for cancer, the central nervous system, and the cardiovascular system, contain some form of fluorine.

Current research

Although fluorine can be toxic when the concentration within the body is too high, it can also be a beneficial element to include in cancer drugs, according to a article published in the Journal of Fluorine Chemistry. According to the research, replacing carbon-hydrogen or carbon-oxygen bonds with a carbon-fluorine bond in the active components of the drug usually shows an improvement of the drugs' effectiveness, including higher metabolic stability, increased binding to target molecules, and enhanced membrane permeability. It is hoped that with the increased effectiveness of the drugs, in conjunction with tumor-specific target drugs or targeted drug delivery systems, the quality of life of cancer patients can be greatly improved over traditional methods such as chemotherapy, in which cancer cells, as well as healthy cells, are targeted by the drugs.

This new generation of cancer-fighting drugs, as well as fluorine-probes to deliver the drugs, has been tested against cancer stem cells and has shown promise in targeting and fighting the cancer stem cells, according to the study. The researchers found that the drugs that included fluorine were several times more active against various cancer stem cells and exhibited better stability than traditional cancer-fighting drugs.

Additional resources

This article was updated on Nov. 24, to include information about the chemical synthesis of fluorine and the discovery of fluorine in nature. 

Since high school, Rachel Ross has been looking up toward the stars to understand how the universe works. She has an undergraduate degree from the University of California Davis and a master's degree in astronomy from James Cook University. Rachel has spent several years making her passion for astronomy and science education into a profession. She has even held the position of Jedi master at an observatory. And no matter what anybody says, the final answer is always 42 and duct tape is useful in all situations.
Sours: https://www.livescience.com/fluorine.html

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