fluorine paper

[budullewraagh]

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.

 

i'll post the paper later when i finish

[ed84c]

This may help;;

 

Fluorine

Fluorine, F2, is a poisonous pale yellow gaseous element found in Group VIIb (i.e. the Halogen Group of elements) of the periodic table.

 

* Atomic Number : 9

* Atomic Mass : 18.9984

* Melting Point : -219.62 degC

* Boiling Point : -188.1 degC

* Density : 1.7

 

The liquid has a clear yellow colour.

 

Fluorine is a highly dangerous element, causing severe chemical burns on contact with the skin.

Discovery

 

* Fluorine was first identified by Scheele in 1771AD.

* Fluorine was first isolated by Moissan in 1886AD, when he noted the inclusion of fluorine in crystals of fluorspar from Quincie.

 

Occurrence

Fluorine is an element which widely distributed, but not very abundant. However, because of its high reactivity is not found naturally in its elemental state.

 

* Fluorine occurs in Cryolite, Na3AlF6,

* Fluorspar, CaF2, and

* Fluorapatite, 3Ca3(PO4)2.Ca(ClF).

 

Extraction

Preparation

Fluorine can be prepared by the electrolysis of fused acid potassium fluoride at 373 degK in a copper vessel using graphite electrodes. The electrolyte, pure and dry acid potassium fluoride is fused in an electrically heated vessel, such as a wide copper V-tube, the electrodes are pure acheson graphite rods insulated in bakelite cement stoppers. A current of 5 amperes at 12 volts is used. The fluorine is purified by passage through two copper U-tubes, containing dry sodium fluoride.

Manufacture

Fluorine is made on the industrial scale by the electrolysis of fused acid potassium fluoride in a copper vessel using graphite electrodes at currents up to 2,000 amperes. An anode current efficiency of about 95% was attained. The fluorine gas is then piped in steel or copper vessels and compressed up to 400 lb/in2 in nickel and steel cylinders.

Properties

 

* Fluorine is a light canary-yellow gas at room temperature.

* Fluorine is the most reactive and electronegative of all elements.

 

Reactions

 

* Fluorine reacts explosively with hydrogen even in the dark and at low temperature.

* Fluorine reacts violently with water forming hydrogen fluoride, and liberates oxygen which is highly charged with ozone.

2 F2 + 2 H2O ==> 4 HF + O2

 

* Fluorine also reacts with sulphur, selenium, and tellurium, which melt and ignite in the gas, forming halides.

* Fluorine is a powerful oxidising agent. For example, when fluorine is bubbles through a solution of potassium chlorate this is oxidised to potassium perchlorate.

F2 + 2 KClO3 + H2O ==> 2 HF + KClO4

* Fluorine does not react directly with oxygen, or nitrogen, and combines with chlorine, only on heating, forming the gaseous products chlorine fluoride, and chlorine trifluoride.

* Fluorine readily combines with bromine, and iodine, forming colourless liquids BrF3 and IF5.

 

Uses

Fluorine is used in the synthesis of organic fluorine compounds.

 

Fluorine is used in the manufacture of Freon (i.e. dichlorodifluoromethane, CCl2F2) which is used as a refrigerant.

 

Fluorine is used in the manufacture of uranium hexafluoride which is required for the separation of the isotopes of uranium in centrifuges.

Detection and Analysis

Fluorine is detected by its corrosive action on glass (i.e. the glass is etched).

[Sayonara]

provide a link to

[budullewraagh]

thanks ed but i already have all that:\

i researched for hours.

 

sayo......im sorry man, i looked and couldnt find

[Sayonara]

Huh?

 

I was pointing that out to ed, who felt okay with posting the info but didn't credit his source.

[budullewraagh]

ah ok thats cool then. i thought you were accusing me of being lazy for not finding links myself. in truth, i researched on and offline for hours but there were things i just couldnt find

[Gilded]

http://www.lateralscience.co.uk/Fluorine/Fluorine.html <- some info about the experiment (and fluorine in general)

 

The date-problem is quite odd though. I can't seem to find it anywhere.

[budullewraagh]

yeah i used that site:\

thanks tho

[budullewraagh]

it is finished. i would appreciate feedback if possible.

if there is anything i could possibly add, let me know.

without further adieu, i present my fluorine paper:

 

 

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

| | |

O-Si-O-Si-O-Si-O

| | |

O O O

| | |

O-Si-O-Si-O-Si-O

| | |

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.

[jdurg]

Very nice. The only thing that came to my attention was where the origin of the name was defined. You are correct in that the origin of Fluorine's name is the Latin word 'fluere', but you're incorrect on what 'fluere' means. Fluere means 'To Flow', not 'To Flux'. (The word 'Flux' is also from the Latin word 'fluere' since a flux allows a metal ore to flow better by lowering its melting point).

[budullewraagh]

actually, as mentioned in my paper, calcium fluoride was used to lower the melting point of metal ores. fluorine was first used (as calcium fluoride) in this way. i have sources and it turns out that fluorine comes from the "flux" meaning of "fluere". thanks for your input though

[jdurg]

actually, as mentioned in my paper, calcium fluoride was used to lower the melting point of metal ores. fluorine was first used (as calcium fluoride) in this way. i have sources and it turns out that fluorine comes from the "flux" meaning of "fluere". thanks for your input though

 

 

Okay, but I just wanted to make sure you didn't get docked for little things. (I have always read and always seen many sources where flux was the name given to CaF2 since it was used to make the metal ore 'flow' better. Also, the word 'flux' is derived from the latin word 'fluere').

[budullewraagh]

yeah no worries. a ctually this isn't getting graded. in fact it's not for a class. speaking of classes, im not even taking chem this year. i'm taking AP physics and AP bio and i'll be taking the AP chem EXAM at the end of this year and the actual class next year. i wrote this paper for apps for summer college

[Silencer]

lol... taking the exam before the class! more power to you.

 

i mean, people in my class think i'm really smart (and a major nerd) because i knew most of the chemistry we learned. I can only imagine how you are viewed by your peers.

[jdurg]

What's funny is that when you're young, people tend to snicker and laugh if you are into chemistry and know a lot about it. Yet when you get older, I've found that people become more impressed and whatnot. For you budullewraagh, the AP Chemistry test should be know problem. I was able to score a 5 on it without any problems at all. The SAT II in Chemistry is another good exam to take. It's graded like the SATs, but on a scale of 200-800. (You get 200 for putting your name on the paper). I took that test twice and got a 790 the first time, but took it again because I knew I could do better. On the second try I scored an 800. For some reason, I always do very well on standardized chemistry exams. I had to take the standardized ACS exams for my organic and physical chemistry courses, and if not for those exams I may not have passed. However, I did so well on them that my grades for the courses were boosted tremendously.

[budullewraagh]

i mean, people in my class think i'm really smart (and a major nerd) because i knew most of the chemistry we learned. I can only imagine how you are viewed by your peers.

i get respect from my peers

my former chem teacher think i'm insane though and is somewhat afraid of me :\

 

 

is there ANYTHING i could add to my paper? is there ANYTHING i should change?

[jdurg]

lol. I know that feeling. My high school chemistry teacher was great. She let me come into the lab during my study hall and just go through the volumes upon volumes of chemistry demonstration books that she had and test some of them out. She also gave me free roam of the chemical supply closet. (There was some scary stuff in there. A half gallon bottle of bromine, about a quarter-pound lump of white phosphorus, some slabs of potassium and sodium, a canister of chlorine, etc. etc. Heh. Too bad the high school couldn't donate some of that stuff to me. ) She also gave me the kindest, most helpful reference letter when I applied to various colleges. I spoke with her via e-mail about a year ago and mentioned my element collection. She asked me to come down to the school one day to show off some of my pieces. So I brought out my cesium, osmium, indium, phosphorus, tungsten, and barium samples. It felt kind of good being up in front of a classroom and having everybody paying close attention to my every word. (This was an AP Chemistry class, so the kids in there actually wanted to be in there). She also downloaded my periodic table photograph and had it printed out on a giant poster sheet. That was pretty cool to see. I should probably drop her a line and see how things are going.

[budullewraagh]

yeah, while the honours chem teacher is afraid of me, the ap chem teacher is so cool...perhaps thats because when i met her i asked to borrow the ap chem book so i could study for the exam. finished through page 200 the day i got it

[Gilded]

"forming hydrofluoric acid and krypton gas or, depending on pressure, K8(H2O)46, a clathrate complex of krypton."

 

Isn't it supposed to be Kr8(H2O)46?

 

And good to see that you have/had good chem teachers. I had a good one at 8th-9th grade, but it seems like I'm going to have a rather strict one in the future (the type that goes "EEEK!" if you try an experiment involving diprotic acids and pine tables).