All About
Elements:
Fluorine
1
Ward’s All About Elements Series
Building Real-World Connections to
the Building Blocks of Chemistry
PERIODIC TABLE OF THE ELEMENTS
GROUP
1/IA
18/VIIIA
1
H
1.01
35
2/IIA
3
Li Be
6.94
13/IIIA
Symbol
22.99
10.81
Atomic Weight
13
24.31
19
3/IIIB
20
4/IVB
21
5/VB
22
Ca Sc Ti
39.10
40.08
37
44.96
38
Rb Sr
85.47
Y
87.62
55
56
39
57–71
´
Ac
- Lr
Ac-Lr
Fr Ra
´´
(226.03)
(226)
40
6/VIB
7/VIIB
24
9
VIIIB
VIII
8
25
26
10
27
11/IB
28
29
12/IIB
30
15/VA
6
C
12.01
14
7
N
14.01
15
Al
Si
P
26.98
28.09
30.97
31
32
16/VIA
33
8
O
16.00
S
16
32.07
72
73
Hf Ta
180.95
104
105
19.00
17
10
Ne
20.18
18
Cl
Ar
35.45
39.95
34
35
36
Cr Mn Fe Co Ni Cu Zn Ga Ge As Se Br Kr
52.00
41
92.91
9
F
2
4.00
17/VIIA
54.94
42
55.85
43
44
58.93
58.69
45
63.55
46
47
65.41
48
69.72
49
72.64
50
74.92
51
78.96
78.96
52
Zr Nb Mo Tc Ru Rh Pd Ag Cd In Sn Sb Te
178.49
88 89–103
23
50.94
91.22
La-Lu
137.33
87
V
47.87
88.91
Cs Ba
132.91
14/IVA
5
B
79.90
12
Na Mg
(223.02)
(223)
Atomic Number
Br
4
9.01
11
K
He
KEY
95.94
95.94
(98)
(97.91)
74
101.07
75
76
W Re Os
183.84
106
186.21
107
190.23
108
102.91
77
Ir
192.22
106.42
107.87
78
79
112.41
80
114.82
81
118.71
82
121.76
83
127.60
84
79.90
I
53
126.90
85
83.80
54
Xe
131.29
86
Pt Au Hg Tl Pb Bi Po At Rn
195.08
109
196.97
110
111
200.59
112
204.38
113
207.2
207.20
114
208.98
115
(208.98)
(209)
116
(209.99)
(210)
117
(222.02)
(222)
118
Rf Db Sg Bh Hs Mt Ds Rg Uub Uut Uuq Uup Uuh Uus Uuo
(261)
(261.11)
´
(262)
(262.11)
57
(266)
(266.12)
58
(264)
(264.12)
59
(277)
(277.00)
60
61
(268)
(268.14)
(269)
(247.07)
62
63
(272)
(280.00)
64
(285)
(285.00)
65
(284)
(284.00)
66
(289)
(289.00)
67
(288)
(288.00)
68
(289)
(293.00)
69
(294.00)
70
(294)
(294.00)
71
La Ce Pr Nd Pm Sm Eu Gd Tb Dy Ho Er Tm Yb Lu
138.91
US: www.wardsci.com
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´´
© 2010 Rev. 7/15 Ward’s Science. All Rights Reserved.
No portion of this work may be reproduced in any form
or by any means without express prior written permission
from Ward’s Science.
US: www.scholarchemistry.com
Canada: www.scholarchemistry.ca
866-260-0501
89
© Copyright 2010 ScholAR Chemistry. All Rights Reserved.
140.91
90
91
Ac Th Pa
(227)
(227.03)
NOTES:
BASIC
PERIODIC
- SIDE 1 (Rev.
7/15)
No portion
of this workTABLE
may be reproduced
in any form
or by any means
140.12
232.04
232.04
Black — solid
Red — gas
231.04
Blue — liquid
144.24
U
92
238.03
(145)
(144.91)
93
150.36
151.97
151.96
94
157.25
95
96
158.93
97
162.50
98
164.93
99
167.26
100
168.93
101
173.04
102
174.97
103
Np Pu Am Cm Bk Cf Es Fm Md No Lr
(237)
(237.05)
(244)
(244.06)
(243)
(243.06)
(247)
(247.07)
(247)
(247.07)
(251)
(251.08)
(252)
(252.08)
(257)
(257.10)
(258)
(258.10)
(259)
(259.10)
(262)
(262.11)
— synthetically prepared
Values provided are based on the 85th edition of the CRC Handbook of Chemistry and Physics. Some values have been rounded.
without express prior written permission from ScholAR Chemistry.
Catalog #9630200
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2
Fun Facts
About…
Fluorine
1. F luorine has both toxic and helpful properties that make it widely used, from medical
treatments to poisonous gases and nuclear
applications.
2. F luorine gets its name from the Latin word
“fleure” which means “to flow”, which is
what fluorine allows metals to do. 1
3. A
lthough the element fluorine was known
to exist in 1813, it wasn’t until 1886 that
Henri Moissan was able to isolate the element after numerous failed, and sickening, attempts.
F
9
19.00
4. F luorine wasn’t mass produced until it was needed for nuclear applications during World
War II.
5. O
ne of the reasons you’ve probably had only a cavity or two is due to the fluoridation of
water, which was first tested in 1944 in Grand Rapids Michigan. Here it was determined that
fluoride helped strengthen teeth.
All About Fluorine:
Fluorine is the ninth element on the periodic table
found in group 17 (7A) with an atomic number of 9 and
a symbol F. Fluorine gets its name from the Latin word
“fleure” which means “to flow”, which is what fluorine
allowed metals to do. 1 Fluorine came also from the “fluor”
in fluorspar (which we now know to be calcium fluoride.)1
Group 17 is known as the halogens and fluorine is known
to be the most reactive of all of the halogens, gaining one
electron in most cases to become a stable fluoride ion.
Fluorine is the smallest of the halogens with an electron
configuration of 1s22s22p5, and has both the highest
ionization energy in its group and the highest electronegativity on the periodic table. Fluorine has a density of
1.696 grams per liter which is the lowest density of all of the elements in its group and is the second
least dense gas, behind neon with a density of 0.9 g/L, in its period on the periodic table. 4
Before fluorine was isolated in 1886, it was known to exist in a number of compounds.1 These
compounds are where you will find fluorine most of the time today, as diatomic elemental fluorine
doesn’t exist in nature, but rather must be created by electrolysis. Some of the more common
compounds of fluorine are fluorspar, hydrofluoric acid, potassium fluoride, sodium fluoride, tin
(II) fluoride, sodium monofluorophosphate, uranium hexafluoride, dichlorodifluoromethane, and
polytetrafluoroethylene.5 These compounds range in use from additives to toothpaste for strengthening teeth, to use in atomic bombs for their explosive nature. Fluorine has also been found to react
with the noble gases xenon, krypton and radon despite their full octet of electrons. 6 Although it is
found in these and other compounds, fluorine typically only has an oxidation state of -1 due to its
high electronegativity. In addition, there is only one stable isotope of fluorine, F-19 with 9 protons
and 10 neutrons.5
3
Properties of Fluorine
In its pure gaseous state, fluorine is pale yellow and extremely
corrosive reacting with most organic and inorganic substances.6
Specifically, fluorine’s color can only be seen by looking down at
it through a test tube, not from the side.18 In addition, it can be
lethal and will react with glass, ceramics, carbon, finely divided
metals and even water to create an explosion.6 Because of its
explosive nature, it is commonly shipped as a cryogenic liquid.10
If fluorine is inhaled or absorbed through the skin it is toxic, but
even at lower than lethal concentrations it can cause chemical
burns.10 Some of these properties have been put to good use,
as hydrofluoric acid is used to etch glass, something that is a
delicate process.
Even though fluorine is found in trace amounts on Earth (it is the 13th most common element in
Earth’s crust,11 and only amounts to 1.3 milligrams per liter in the ocean5) and in the universe, it is most
commonly found in the form of a fluoride ion. For instance, the fluoride ion is present at about 3 ppm
in the human body!7 Compounds that are added to toothpaste and drinking water are all in the form
of the ion, not the toxic gas. All of the stable naturally occurring fluorine is F-19, with isotopes used
for a variety of medical purposes and having a variety of half lives. The half lives range from 4.557 x
10-22 seconds for F-15 to 109.77 minutes for F-18 making the radioisotopes quickly removed from the
body.9 This isotope with 9 protons and neutrons, F-18, is commonly used as a radioisotope to detect
cancer, for cardiac and brain imaging as well as monitoring the progress of treatment of these diseases
because it has been “proven to be the most accurate non-invasive method of detecting and evaluating
most cancers.” 8.
Discovery and History
If you thought, based on some of the properties described, that fluorine had a rather interesting history you’d be correct! First, Fluorine is created in stars and quickly breaks down producing
oxygen and helium or neon and hydrogen after undergoing fusion with hydrogen or helium respectively 16. Although it wasn’t known exactly how it worked, the first use of the mineral fluorite
(made from fluorine and calcium) was for flux 12. Flux is introduced when smelting ores and allows
the metal to flow while it removes the impurities, or slag 12. This is part of where the name came
from for fluorine, as fleure means to flow which is what this mineral allowed metals to do. The
first record of fluorine is from 1670 in a set of instructions on how to etch glass with Bohemian
emerald, now known to be CaF2 5. Specifically the chemist Schwandhard found that glass could
be etched when it was exposed to fluorspar treated with acid.
Early 1800’s
In the early 1800s numerous chemists carried out experiments
on fluorspar (now known to be calcium fluoride) including Gay
Lussac, Humphry Davy and Joseph Priestly. In fact, none of them
were able to isolate fluorine, but many produced hydrofluoric acid,
which unfortunately led to fatalities and to blindness.1 Davy then
used the hydrofluoric acid and became ill trying to isolate fluorine.
Davy and a French scientist, Andre-Marie Ampere, exchanged
letters on how the acid may have a new element in it and the element, although not yet isolated, was given its name of fluorine by
Ampere.1 Finally in 1886, a French chemist named Henri Moissan
was able to isolate fluorine using electrolysis of dry potassium
hydrogen fluoride in anhydrous hydrofluoric acid contained in a
Moissan’s fluorine cell, from his
platinum container that he cooled to -31oC.1 Moissan’s work was in- 1887 publication
terrupted four times due to poisoning caused by the fluorine he was
trying to isolate, but his work was awarded in 1906 with the Nobel Prize in chemistry. 1 It wasn’t
until World War II that commercial production of fluorine began to help with nuclear energy and
the bomb project. 6
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4
Fluoridated Water
In addition to elemental fluorine, fluoridation of water and compounds containing fluorine have
interesting histories. Fluoridated water came to be after researchers in 1901 in Colorado Springs
noticed that people had black and brown spots all over their teeth in addition to “mottled tooth
enamel.”13 They traveled the country to other sites where people lived with cracked teeth trying to
determine what might be causing the problem. In 1923 the problem was found to be the water
source, but it wasn’t until 1931 when photospectrographic analysis was performed on samples
of water that the problem was realized.13 It was also known that the mottled tooth enamel was
resistant to decay. Dr. Trendley Dean wondered if fluoride added to drinking water at lower levels
could help fight decay. After a few years of research and in discussions with the city commission
of Grand Rapids Michigan, fluoride was added to the water there in 1944. The tooth decay rate
of over 30,000 children that were studied decreased by 60 percent.13 According to the CDC, most
additives that are used in the US come from phosphorite rock that is treated with sulfuric acid
to produce a phosphoric acid- gypsum slurry.14 Once heated, this slurry releases both hydrogen
fluoride and silicon tetrafluoride gases which are captured and then condensed into water to use
as fluoride ions.14
Where in the world is Fluorine?
Fluorine is only found in Earth’s crust in an abundance of
0.054%2 and in the universe as 0.00004% abundance.3 In
Earth’s crust it can be found in coal, clay and rocks, as well
as in the soil with approximately 330 parts per million
on average. 16 Fluorite, which is the most common form
of fluorine, is found in numerous locations around the
world including Argentina, Canada, China, France, Mexico,
Russia, Spain and the United States.15 That being said, the
primary areas for mining are China, Mexico and western
Europe.16 Depending on where the mineral containing
fluoride is collected, it will fluoresce different colors. Fluorite is one of the more popular minerals
in the world among mineral and gem collections, despite its toxicity when treated and turned
into hydrofluoric acid or fluorine gas.15 Fluoride ions are added to water in the United States so
that the concentration is about 0.7 milligrams per liter and because fluoride can be natural in
some places (coming from minerals) the maximum level allowed in 4.0 milligrams per liter.14
Compounds containing fluorine also exist in the atmosphere, depleting the ozone as CFC’s or
chlorofluorocarbons. These were originally used as propellants in many spray cans are harmless
to humans. Sadly, it was later discovered that the chlorine reacts with ozone, O3, removing the
protective layer from the earth’s atmosphere. Since 1996, CFC’s have been banned and hydrofluorocarbons (HFCs) have been created and used as replacements. These compounds are specifically used in refrigerators, air conditioning systems and as fire retardants. HFCs are believed to be
safer than CFCs for the ozone layer and are more controlled in terms of where they are used.17
Uses of Fluorine Today:
Elemental fluorine is so highly reactive and toxic to humans that
it is rarely used today. It can even react with diamond, one of the
strongest materials on Earth. Compounds containing fluorine
are much more common and useful than diatomic fluorine gas
in its normal state. For example, gem collectors search the globe
to obtain different colors of nephrite, topaz and feldspar each of
which contains fluorine components. Hydrofluoric acid, although
it is a weak acid, is highly toxic to humans, but is used in etching
and frosting glass. Fluorite can be used in the processing of metals and smelting of iron. Another very important use of fluorine
is in refrigerants for both air conditioners and refrigerators that
have be changed since CFS’s were banned to other fluorine
containing compounds. Below you will find other applications of fluorine.
5
Teflon:
Teflon, although commonly thought of today as toxic,
was once used on all non-stick pans. It is produced,
along with other fluoropolymers, around 180,000
metric tons annually.16 Teflon, Polytetrafluoroethene,
is considered a high temperature plastic that doesn’t
melt until 326°C (or 620.3°F for those of us used to
baking at temperatures upward of 400°F). In addition
to being used in cooking, it is also used in cable insulation, plumber’s tape and as the base of Gore-Tech®
which is used to waterproof shoes and clothing. 7
http://home.howstuffworks.com/nonstick-cookware.htm
Nuclear Power and Bombs:
Uranium hexafluoride, UF6, is helpful in both nuclear reactors and bombs in processing nuclear fuel.
UF6 can be used in all three states of matter: solid for storage, liquid for filling or emptying containers or equipment and as a gas for processing.19 The uranium that is used has to be enriched so the
percent of uranium 235 is higher. The UF6 helps this process take place by enriching one sample and
depleting the other sample effectively “transferring” the U-235 from one sample to another. 20
Breaking Bad:
Dissolving Bodies?
If you’ve ever seen the popular T.V. series, Breaking
Bad, you know Walter White and his partner in crime
Jesse Pinkman. In one episode, they tried to dispose
of a rival’s body in a bathtub using hydrofluoric acid
and ran into some trouble. Instead of pouring the
contents into a plastic container (which hydrofluoric
acid CAN’T dissolve), Jesse pours the HF into a bathtub, which is then dissolved by the acid. Later, the
tub and disintegrating body fall through the ceiling
http://www.bbc.com/news/magazine-23710654
onto the floor below. While there is some truth to
this concept, in reality, it is unlikely that the tub would
go through the ceiling in this manner. The episode was even debunked by Mythbusters, which found
that after eight hours of soaking in HF, all the materials that would have been between the person
they were trying to dispose of and the next floor, had deteriorated to different degrees, but none were
completely destroyed 21.
Chemotherapy:
The drug Fluorouracil is used in chemotherapy
treatments for cancers ranging from breast and
colon to cervical and head cancer. It can also be
used as a topical cream on basal cell cancer.22 This
specific chemotherapy drug is considered an
antimetabolite, which is very similar to substances
normally in cells and therefore easily integrated.
Fluorouracil interferes specifically with DNA and
RNA synthesis by “mimicking the building blocks
necessary for synthesis.” 22
6
Other Pharmaceuticals:
In addition to being used in chemotherapy,
which is known for its adverse side effects,
fluorine containing compounds are used in numerous other medicines. The element appears in
approximately 20% of all pharmaceuticals mainly
due to the strength of the bonds it forms with
other atoms—specifically carbon.7 Since fluorine
is so reactive, it isn’t found alone in nature and
the bonds that it forms are strong due to fluorine’s high electronegativity. Many times fluorine
is used to protect the drug from degradation
helping to last longer and extending the active
lifetime inside the patient’s’ body.7 To do this,
scientists replace hydrogen with fluorine creating
a stronger bond between elements. In addition, fluorine can help change the shape of the
molecule it is introduced in helping it bind better to the protein it is targeting.7 Common drugs that
use these properties of fluorine include Prozac, an anti-depressant, and Lipitor, a cholesterol-lowering
drug. 7.
Rocket Fuel?
Elemental fluorine, F2, has been studied as a
possible rocket propellant.6 Engines that run
on liquid fluorine have been developed and
fired successfully, but are not widely used due
to safety concerns.23 For example, fluorine is
extremely toxic; It is considered a “super oxidizer”
and it reacts violently with almost anything
except substances that have already been fluorinated, lighter noble gases and nitrogen.23 Liquid
fluorine has also been mixed with liquid oxygen
to improve the performance of LOX- burning
engines to create FLOX (Fluorine-Liquid-Oxygen
burning). Despite its ability to propel and
explode, due to safety concerns around toxicity,
fluorine has been abandoned by most nations
hoping to travel to space.
Toothpaste:
Pearly whites. Chompers. What have we to thank
for these novel nicknames other than toothpaste.
Fluoride-containing compounds are added to
toothpaste to help strengthen the enamel and
prevent cavities similar to how fluoride is added
to water.24 When toothpaste is applied at the
tooth buds, the hydroxyapatite enamel of the
tooth is replaced by fluorapatite. This compound, fluorapatite, is more resistant to decay
than hydroxyapatite because it is less soluble in
acid than hydroxyapatite 24.
7
References:
1. h
ttp://www.chemicool.com/elements/fluorine.html
2. http://periodictable.com/Properties/A/CrustAbundance.an.html
3. http://periodictable.com/Properties/A/UniverseAbundance.v.html
4. http://www.angstromsciences.com/density-elements-chart
5. h
ttp://education.jlab.org/itselemental/ele009.html
6. h
ttp://www.livescience.com/28779-fluorine.html
7. h
ttp://www.rsc.org/periodic-table/element/9/fluorine
8. h
ttp://www.world-nuclear.org/information-library/non-power-nuclear-applications/
radioisotopes-research/radioisotopes-in-medicine.aspx
9. http://education.jlab.org/itselemental/iso009.html
10. https://cameochemicals.noaa.gov/chemical/764
11. http://healthresearchfunding.org/10-interesting-facts-fluorine/
12. https://www.britannica.com/technology/flux-metallurgy
13. http://www.nidcr.nih.gov/oralhealth/Topics/Fluoride/TheStoryofFluoridation.htm
14. http://www.cdc.gov/fluoridation/factsheets/engineering/wfadditives.htm
15. http://www.gemselect.com/gem-info/fluorite/fluorite-info.php
16. http://www.lenntech.com/periodic/elements/f.htm
17. https://toxtown.nlm.nih.gov/text_version/chemicals.php?id=9
18. http://www.softschools.com/facts/periodic_table/fluorine_facts/179/
19. http://web.ead.anl.gov/uranium/guide/uf6/
20. http://web.ead.anl.gov/uranium/guide/depletedu/enrich/index.cfm
21. http://mythresults.com/breaking-bad-special
22. http://chemocare.com/chemotherapy/drug-info/fluorouracil.aspx
23. http://www.braeunig.us/space/propel.htm
24. http://www.scientificamerican.com/article/how-does-the-fluoride-in/
8
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