budullewraagh
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Discovery of Fluorine
can anyone provide a link to information about moissan's experiment? i know that he used a cooled platinum u-tube, KHF2, HF and iridium-platinum
alloy electrolytes and that he did this in 1886, but i do not know the exact date and i also feel there are a few other things i could incorporate in
a paper i am writing on fluorine. as a sort of prologue i intend to write a paragraph on the actual time and place of discovery of fluorine.
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BromicAcid
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I don't have any links but I did look up quite a bit of the information, much information on his first attempts is available partially because he
won the Nobel Prize for this an his accomplishments with the furnace design. One little side detail, he succeeded in making fluorine in the vessel
you describe, cooled with methyl chloride, however the next day when he brought people over to witness the experiment it failed, it failed because he
was attempting to electrolyze pure HF, in his previous attempt it was contaminated with KF and he was unaware it made a difference, later he added a
small amount of KF and the electrolysis worked. There are three different versions of the electrolyte in the cell, the low temperature version that
he pioneered, a moderate temperature version, and the high temperature version that uses straight potassium bifluoride that operates above 100C. I
have quite a bit more information that I can dig up later if you so wish. There should be a picture of his original cell in the thread on how to make
fluorine, I asked and received two or three of them I believe.
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JustMe
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This is the thread Bromic Acid was referring to:
http://www.sciencemadness.org/talk/viewthread.php?tid=2358
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budullewraagh
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thanks for your help. i needed information for a paper i wrote.
i would appreciate feedback if possible. if there is anything i could possibly add, let me know. the paper is as follows:
One night in the Sarbonne region of France in the year 1886 Henri Moissan nervously prepared his apparatus. Having already spent a great amount of
money on custom made parts, featuring a platinum U-tube, Moissan knew success was necessary to his professional career, if not his life. The skittish
Moissan thought over the procedure of his experiment one last time. This could not fail he thought, a nagging doubt nonetheless in his mind as he
turned on his direct current battery. Soon Moissan noticed a pale yellow gas appear at the anode. Success was indeed apparent after almost a century
of constant experimentation. The tiger had finally been isolated.
Fluorine (Latin, fluere, translation to "flux" always proved to be an
elusive element. First isolated in 1886 by the French chemist Henri Moissan, fluorine was originally referred to as the "gas of Lucifer" and
the great "tiger" of chemistry. Such names were true to its remarkable properties, which is what left fluorine as a mysterious, almost
mythological element for decades, as most chemists chose not to attempt to isolate it for fear of death. In fact, many of the first chemists who
isolated fluorine suffered significant physical harm, some of which proved fatal, as fluorine attacks the heart and bones rapidly. Fluorine has been
mentioned in realistic fictional writings such as The Ernest Glitch Chronicles. In these writings, one can see the mystery this chemical proved to be
in the years following its discovery. In a fictional letter written to Michael Farraday on July 8, 1856, Ernest Glitch wrote "Your concern about
Hodges` health and employment exasperates me Faraday! However, in accordance with your wishes, I have reinstated him. I may say Faraday, that finding
Hodges afforded me no end of trouble, and not inconsiderable expense. The wretch had travelled to Newcastle town moor, and was attempting to eke out
an existence with a carnival show. Hodges` earning potential had recently plummeted, like a stooping peregrine, after an unfortunate incident
involving the mayors wife. The carnival owner forbade him to further display his fluorine scars, and his electrical scarification is hardly of the
shocking nature lucifers gas, the tiger of chemistry, had afforded him. He is still maintaining blindness in one eye, and I have had to reduce his
previous workload because of frailness." Fluorine was always considered to be a remarkable element, even in its salt forms.
In 1529, Georgius Agricola discovered that fluor lapis, (now known as fluorspar or CaF2) with its low melting point of 1403 Celsius could be used as a
flux to lower the activation energy of the melting of ores for smelting. One and a half centuries later in 1676, it was found that flux, when heated
on a metal plate, shines with a blue-white luster.
This would later be found to be a result of catalytic action of the fluoride anion on ozone, which causes colourless diatomic oxygen to oxidize to the
pale blue ozone (O3) molecule. At about the same time, a glass cutter found that heating fluorspar with strong acids created agents that etched glass.
In 1771, Carl Wilhelm Schelle discovered hydrofluoric acid (HF) after heating fluorspar with oil of vitriol (sulfuric acid, H2SO4). Schelle also found
that this acid was capable of dissolving glass and silaceous earth. In discovering this acid he helped Humphry Davy disprove Antoine Lavoisier’s
theory that all acids contain oxygen. André-Marie Ampère proposed to Davy that the acid must contain hydrogen and a new undiscovered element. In the
experimentation the ensued, Davy, Joseph Louis Gay-Lussac and Louis Thenard suffered greatly from inhalation of the acid vapours which predominantly
attack the heart and bones. The Reverend Thomas Knox and his brother George made an apparatus with fluorspar. The Reverend barely escaped with his
life and George was bedridden for three years before he regained his health. Paulin Louyet of Brussels, Belgium nonetheless continued his
experimentation and died, as did Jérôme Nicklès of Nancy, France. Edmond Fremy attempted to replicate Louyet’s experiments but instead of using
hydrated fluorspar he electrolyzed anhydrous fluorspar. Unfortunately, the gas produced at the anode reacted instantly. In 1869 George Gore produced a
small amount of fluorine through electrolysis. Unfortunately for Gore, he had not realized that fluorine would react explosively when combined with
hydrogen and the two gases produced at opposite electrodes exploded violently. Gore narrowly escaped injury. Fluorine was not successfully isolated
until 1886 when Ferdinand-Frédéric-Henri Moissan, a student of Fremy’s successfully electrolyzed dry potassium acid fluoride in anhydrous liquid
hydrogen fluoride in a platinum U-tube using platinum-iridium electrodes using methyl chloride to cool the apparatus to -50 Celsius. Although an
advocate of extreme cleanliness in his laboratory, Moissan was subjected to fluorine poisoning many times and died prematurely as a result. For one to
understand the dangers and true beauty of fluorine, one must consider its remarkable properties.
Fluorine, element number 9, has the greatest electronegativity of all elements according to the Pauling scale as well as the Allred-Rochow scale.
According to Linus Pauling, electronegativity can be defined as the relative attraction of an atom for the valence electrons in a covalent bond. It is
proportional to the effective nuclear charge and inversely proportional to the covalent radius. Pauling’s electronegativity formula is as follows:
X=(0.31(n+1+or-c))/r+0.50
where the variable "n" represents the number of valence electrons, "c" represents any formal valence charge on the atom, where the
sign before it corresponds to the sign of this charge and "r" represents the covalent radius in angstroms. One can easily calculate the
electronegativity of fluorine using Pauling’s formula:
X(F)=(0.31(7+1-1))/0.64+0.50
X(F)=3.890625
Pauling realized that because electronegativity is not concerned with atoms in isolation, but rather atoms in molecules, it is not possible to define
precise electronegativity values. All of Pauling’s figures were based on bond energy data measured using heats of dissociation and formation. It was
because of the lack of precision of electronegativity values that Pauling originally gave fluorine an arbitrary electronegativity of 4.0. For cations,
however, this formula has proven to be less than accurate. A new formula was proposed for the electronegativity of cations:
X=0.44(phi)-0.15
where (phi) represented the work function, or the energy necessary to remove an electron from the metal surface in thermoelectric or photoelectric
emissions. According to this formula, the electronegativity of calcium is equal to 1. According to Pauling’s formula,
X(Ca)=0.31(2+1+2))/1.74+0.5
X(Ca)=1.390804598
Now, considering fluorine’s electronegativity, one can understand how fluorspar is able to catalyze the oxidation of diatomic oxygen to ozone.
According to the Gordy formula, the degree of ionic character is equal to the absolute value of the difference of the electronegativities of bonded
atoms divided by 2:
DIC=|Xa-Xb|/2
DIC=|3.890625-1.390804598|/2
DIC=1.249910201
More accurately:
DIC=|3.890625-1|/2
DIC=1.4453125
A more sophisticated formula was proposed by Hannay-Smyth:
DIC=0.46*(|Xa-Xb|)+0.035*(Xa-Xb)^2
DIC=0.46*(|3.890625-1|)+0.035*(3.890625-1)^2
DIC=4.254187012
Calcium fluoride proves to be very ionic, and at high temperatures proves to be even more ionic due to the increased rate of collisions between
molecules.
The bond stretching force constant for stable molecules exhibiting their normal covalencies can be estimated (in units of 10^5 dynes per cm^-1) by the
formula:
K=1.67N((XaXb)/D^2)^(3/4)+0.3
where "N" represents the bond order and "D" represents the internuclear distance in angstroms.
K=1.67*1((3.980625-1)/(.64+1.74)^2)^(3/4)+0.3
K=1.331766779
These extreme numbers are simply numerical representations of the physics of the fluorine atom. It is the combination of these physical properties
that gives fluorine its characteristic bonding abilities. Fluorine is atomic number nine and has a valence shell containing two electrons in its 2S
shell and five electrons in its 2P orbital. Because its valence shell is only one electron short of satisfying the octet rule, there is the maximum
possible number of protons present before the electrons must be placed significantly farther away in a 3S orbital (disregarding the inert gas neon).
With covalent radius of 64 picometers, fluorine is quite small compared to other halogens. The radii (in picometers) of chlorine, bromine and iodine,
respectively, in a diatomic molecule are 99, 114.2 and 133.3. Despite the fact that chlorine has more protons than fluorine, bromine has more protons
than chlorine and iodine has more protons than bromine and the fact that the effective nuclear charges of the halogens increases down group 17, the
protons lose their attractive force and control over the increased distance from the nucleus to the valence shell, leaving fluorine as the element
with the greatest forces to attract electrons. This property is clearly evident in its reduction potentials which are, (in E^(theta)/V):
F2 __2.866__ F-(aq)
F2 __2.979__ HF2-
F2 __3.053__ HF(aq)
as well as its ionization energies (in kJ/mol^-1) the first of which is 1,681.
These atomic and valence properties give fluorine its remarkable bonding properties. Such properties made fluorine remarkably difficult to isolate
during the early years of its experimentation. When pumped into an ionic aqueous solution, fluorine will displace any anion, sometimes resulting in
decomposition of this anion. Examples can be demonstrated by the following reactions:
F2(g)+2NaCl(aq)à2NaF(aq)+Cl2(g)
F2(g)+2KMnO4(aq)à2KF(aq)+2MnO2(s)+2O2(g)
While completely dry fluorine cannot attack completely pure amorphous silicon dioxide, it is almost impossible to completely dry fluorine and
eliminate impurities in glass. As a result, there are multiple ways in which fluorine is able to self-catalyze. If the fluorine is wet at all, it
hydrolyzes somewhat to form hydrogen fluoride, which attacks glass. If the fluorine is dry, it finds impurities in the glass, which ideally has the
tetrahedral structure:
O O O
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O-Si-O-Si-O-Si-O
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O O O
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O-Si-O-Si-O-Si-O
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O O O
Of course, this is not a possible structure, as all of the oxygen atoms bordering the outside of the structure are lacking one bond. The above would
be more similar to the quartz lattice structure of silicon dioxide than glass. Glass itself is an amorphous solid, which contains various impurities.
One example of an impurity is SiO(OH)2. While dry fluorine is not capable of attacking a double bond between silicon and oxygen, it is able to attack
the hydroxyl radicals which are repelled by the doubly bonded oxygen, which is situated 1.7595 angstroms away form the silicon atom. Because it is so
far away from the silicon, the doubly bonded oxygen does not effectively repel fluorine as it attacks the silicon and hydroxyl radicals. Diatomic
fluorine immediately forms hydrogen fluoride on contact with the hydroxyl radicals. The oxygen of the hydroxyl radicals, having lost a bond,
instantaneously bonds together and is displaced by the hydrogen fluoride, forming SiOF2H2 as an intermediate and O2 which is immediately attacked by
more fluorine to form gaseous tetrahedral SiF4 and H2O. The water produced hydrolyzes fluorine to form yet more HF, which directly attacks even pure
silicon dioxide. In this reaction, hydrogen bonds to oxygen and fluorine bonds to silicon (4HF+SiO2àSiF4+2H2O). Other products and intermediates of
these reactions are tetrafluorosilicic acid and hexafluorosilicic acid.
The ability of fluorine to attack many oxidized compounds has proven to be quite useful in the way it is able to amplify certain properties of some
compounds. While sulfuric acid (H2SO4) is a very strong acid (pKa1=-3, pKa2=1.99) fluorosulfonic acid (HFSO3) is approximately 10,000 times stronger
(pKa=-8.60). The ability of fluorine to increase dissociation in polyatomic anions can be best demonstrated by pentafluoroantimonic acid (HF:SbF5).
Hydrofluoric acid alone is considered a weak acid, with a pKa of 3.20. When mixed in a solution with antimony pentafluoride however, the HF
dissociates almost completely to form H+ and SbF6-. In this situation, the fluorine completely takes control of the hydrogen’s only atom and forms a
coordinate covalent bond with the antimony pentafluoride, leaving the hydronium cation behind. When fluorine makes a covalent bond, it is likely to
keep this bond and as a result, the equilibrium of this reaction almost exclusively favors the forward reaction. Toothpaste companies have capitalized
on this property. While analyzing the enamel of elephants’ teeth in 1802, Domenico Pini Morichini found much calcium fluoride along with the
predicted predominant presence of calcium phosphate. Bones are composed primarily of hydroxyapatite (Ca2(PO4)3(OH)). When in contact with the fluoride
anion (in the form of aqueous sodium fluoride), the hydroxyl radical is displaced by fluoride which makes an even less soluble compound known as
fluoroapatite. By fluorinating one’s teeth, one can prevent the corrosion of enamel, which is normally slightly susceptible to attack by strong
acids.
Chlorofluorocarbons (CFCs) were synthesized for use in refrigeration devices. These compounds were formed of one to two carbon atoms bonded with
fluorine and some chlorine. While almost completely chemically inert, CFCs were found to be quite hazardous for the atmosphere if released.
Ultraviolet rays from the sun were able to remove the chlorine from CFCs, one molecule of which could catalyze the decomposition of millions of ozone
molecules in seconds. The fluorine however, is bonded so strongly to carbon that it cannot be removed by such rays.
Another strong covalent fluorine compound is polytetrafluorethylene (PTFE), which has the formula (C2F4)n. The carbon in tetrafluoroethylene is SP^2
hybridized, with 120 degree bond angles. In polymerized tetrafluoroethylene, the carbon becomes SP^3 hybridized, forming 109.5 degree bond angles.
With perfectly uniform highly negative fluorine strongly bonded to a long carbon backbone, polytetrafluoroethylene is not only practically chemically
inert, but it also has a remarkably low Mu coefficient of friction with other PTFE, which happens to be 0.04. Considering the fact that friction is
not actual contact between matter, but rather the conflict between repelling forces, the fluorine atoms of PTFE do not touch one another, and are
rather passed, one slightly apart from the other, as they move, with the perfectly uniform PTFE molecules only forming slight "bumps"
between one another in a perfectly symmetrical structure.
Fluorine does indeed make strong covalent bonds, but it also makes somewhat weak covalent bonds with the lower "inert" group 18 gases. Such
compounds as krypton difluoride are formed by pressurized fluorine and krypton gases. This compound hydrolyzes in water, forming hydrofluoric acid
and krypton gas or, depending on pressure, K8(H2O)46, a clathrate complex of krypton. Much more experimentation of the action of fluorine on xenon
has been documented. Such compounds made from xenon and fluorine are XeF2, [XeF]+[AsF6]-, XeF4, XeOF4, XeO2F2, XeF6 and XeO3F2. All fluorinated
xenon compounds are strong oxidizing agents and most hydrolyze to form hydrofluoric acid or a lower xenon fluoride as well as xenon (and xenon
clathrate complexes depending on pressure), and/or the corresponding xenon oxide.
While fluorine tends to make strong covalent bonds, it also forms some unstable ionic bonds. Fluorine nitrate and permanganyl fluoride are incredibly
unstable almost mythical compounds used as rocket propellants. There is almost no information on the latter, but the former has been reasonably well
documented. Formed from the very negative nitrate anion and fluorine, the most negative element, the union of these two compounds is broken remarkably
easily and it decomposes to form diatomic oxygen, nitrogen dioxide and fluorine gas.
Fluorine, the most negative element in existence and the most negative element conceivable, truly lives up to its title of "gas of Lucifer".
With self-catalytic properties, the ability to displace any aqueous anion and the ability to form particularly strong covalent bonds, fluorine is
unique and capable of amplifying the properties of many compounds. These properties have fascinated many and resulted in the advancement of society.
[Edited on 19-1-2005 by budullewraagh]
[Edited on 19-1-2005 by budullewraagh]
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JohnWW
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The isolation of fluorine probably inspired the discovery of the cryolite (Na3AlF6) electrolysis method for producing metallic aluminium from refined
alumina (from bauxite) the same year, 1886. The Al2O3 is dissolved in molten cryolite, which on electrolysis produces Al and F2 which latter reacts
with more Al2O3 to form more cryolite, and so on. This is what is done, for example, at Al smelting plants at Bluff, New Zealand, and in Tasmania,
using low-cost local electricity, and bauxite from Weipa, Queensland, which is refined at Gladstone, Queensland.
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