Sciencemadness Discussion Board - Chemical Isolation of Fluorine

Chemical Isolation of Fluorine

Xenos - 29-8-2002 at 14:44

I know I should probably reply to more existing posts first, but i have this knawing question. I read that in the 80s someone found a way to isolate Fluorine by chemical instead of electrical means. I dont actually plan on doing this! It would most definatly result in death. I was just wondering if anyone knew how this could be done.

Patent: Pure fluorine gas generator

Rhadon - 29-8-2002 at 18:10

US Patent 4,711,680 shows something that could be from interest. Enter the patent number into this form.

Here are some examples
More details can be found in the patent

EXAMPLE 1
A mixture of K₂NiF₆ (0.369 g) and BiF₅ (1.372 g) was rapidly heated, as described above. When the inside temperature reached about 60 °C., rapid fluorine evolution started, resulting in a maximum pressure of 990 torr at a reactor temperature of 170 °C. The purity of the evolved fluorine (1.1 mmol) was shown by mercury analysis to be in excess of 99%.

EXAMPLE 2
A mixture of Cs₂CuF₆ (0.89 g) and BiF₅ (1.20 g) was rapidly heated, as described in Example 1. Again, pure fluorine (0.9 mmol) was evolved, resulting in a maximum pressure of 836 torr.

EXAMPLE 3
A mixture of Cs₂MnF₆ (2.115 g) and BiF₅ (4.515 g) was rapidly heated, as described in Example 1. Again, pure fluorine (1.0 mmol) was evolved, resulting in a maximum pressure of 929 torr.

EXAMPLE 4
A mixture of K₂NiF₆ (1.584 g) and TiF₄ (0.774 g) was rapidly heated, as described in Example 1. Again, pure fluorine (0.87 mmol) was evolved in the temperature range 65 ° to 170 °C., resulting in a maximum pressure of 810 torr.

EXAMPLE 5
A mixture of K₂NiF₆ (0.486 g), TiF₄ (0.240 g) and BiF₅ (0.590 g) was rapidly heated, as described in Example 1. Again, pure fluorine (0.88 mmol) was evolved in the temperature range 60 ° to 180 °C., resulting in a maximum pressure of 820 torr.

Other Hypothetical Methods

PrimoPyro - 29-8-2002 at 19:35

Cheat and use a hypofluorite (if you can find one)

NaOF + 2HF --> NaF + F2 + H2O

If you cannot use that, then oxidize HF to F2 with H2O2.

2NaF + H2SO4 --> Na2SO4 + 2HF
2HF + H2O2 --> F2 + 2H2O

Or like I joked once before, bombard a fluoride salt with alpha particles. Fluorine will not take an electron from a positive sodium ion Na+ because the system of 2Na+ + F2 is more stable than the system of 2Na+2 + 2F-.

2NaF + He+2 --> He + 2NaF+ --> He + 2Na+ + F2

But that isn't realistic, and quite frankly it's cheating isn't it? lol.

PrimoPyro

madscientist - 29-8-2002 at 21:10

Er, you can't oxidize HF to liberate F2. Fluorine is, after all, the most electronegative element. Here's an example of what happens:

3H2O + 3F2 ----> 6HF + O3

Damn, Fluorine Wins Again

PrimoPyro - 30-8-2002 at 13:16

I'm not used to the stubbornness of fluorine; that's twice I've been wrong about it.

I usually consider electronegativities secondly, which has dire consequences with inorganic chemistry where they are very important.

I was thinking with the old basic rule that a reaction usually proceeds if it a)produces water as a product, or b)produces a gas as a product.

PrimoPyro

vulture - 31-8-2002 at 05:35

Heating AgF2?

Rhadon - 31-8-2002 at 06:30

"Chemical Engineering News" (15.09.1986) does also state a method for preparing elemental fluorine by chemical means.

Has anyone acces to this article?

madscientist - 31-8-2002 at 09:48

Heat PbF4 to about 100C and it should liberate F2.

Xenos - 31-8-2002 at 12:07

So why would heating PbF4 or AgF2 free the Fluorine? There has to be a scientific reason.

vulture - 1-9-2002 at 11:33

Well, halogens of silver are notoriously unstable and decompose already under the influence of light, which is used in photography.

PbF4 could also liberate F2 because Pb +4 easily reduces to Pb+2. That's why PbO2 is a relatively strong oxidizer.

a_bab - 20-9-2002 at 10:34

So why the old chemists didn't use this method ? There were lots of intoxicated and even some victims in the race for fluorine isolation as you know.

Nah, I don't believe that it'll work.

Polverone - 20-9-2002 at 11:04

No, I think the method will work. But I don't think you will be able to produce AgF2 without using elemental fluorine at some point. So this is really a method of chemically storing fluorine, not producing it.

BromicAcid - 24-7-2003 at 19:01

[The combustion of compositions containing an excess of oxidizer may involve the evolution of free fluorine in certain cases. The formation of free fluorine will be most likely when the oxidizers used are fluorides of variable valence metals CoF3, MnF3, PbF4, etc. for example:
6CoF3 + 2Mg ---> 2MgF2 + 6CoF2 + F2]
Principles of Pyrotechnics [A.A.Shidlovskiy] Third edition
Maybe this might work with teflon, doesn't fit the description and the reaction would be hard to control but it is more readily availible as a pyrotechnic supply.

Or maybe Xenon Difluoride, it can be used as a fluorinating agent easily and Xenon gas is the only by product, this site has a list of people that sell it http://www.xactix.com/suppliers.htm does anyone know of a place that sells fluorine cylinders?

blip - 24-7-2003 at 23:37

I remember reading about Sir Humphrey Davy decompsing some flourine manganese complex or something to get it, but that was before I knew much about chemistry.

Quote:

NaOF + 2HF --> NaF + F2 + H2O

Fluorine likes to form FOF.
NaOF + HF <strike>  </strike>> FOF + NaOH
NaOH + HF <strike>  </strike>> NaF + H2O

Another example:
H2O + F2 <strike>  </strike>> HF + HOF
HOF + F2 <strike>  </strike>> OF2 + HF

I am a fish - 26-7-2003 at 05:33

The classic method of isolating fluorine chemically is to heat a mixture of antimony pentafluoride and potassium hexafluoromanganate(IV) at 150C. This was discovered in the late 19th century.

BromicAcid - 11-10-2004 at 20:04

A little more detail on I am a Fish's fluorine production method he brought up from:

http://www3.district125.k12.il.us/chemmatt/86/86_12_t/861213...

Quote:

In a short article to be published in an October issue of Inorganic Chemistry, Karl Criste of Rocketdyne Division of Rockwell International describes a “chemical synthesis of elemental fluorine.” Christe first prepares potassium hexafluoromanganate (IV) and antimony pentafluoride by separate nonelectrolytic methods from readily available starting materials:

2KMnO4 + 2KF + 10HF + 3H2O2 ---> 2K2MnF6 + 8H2O + 3O2

SbCl5 + 5HF ---> SbF5 + 5HCl

When these two reagents are heated together at 150 °C for 1 hour, nonelectrolytic “chemical” fluorine is readily obtained.

2K2MnF6 + 4SbF5 ---> 4KSbF6 + 2MnF3 + F2

So many hazardous chemicals.....

Theoretic - 13-10-2004 at 13:26

My Soviet chemistry book (written in 1935, I have the 1965 thirteenth reedition, the author got a prize in 1935 for the coolest chemistry textbook :cool:

) states that MnF3 makes fluorine on heating above 250 C. Also it tantalizingly states that in the prezence of excess HF, KMnO4 is reduced not to the II, but just to the III oxidation state. Could that be the holy grail - KMnO4, HF, reduce with... I don't know, oxalic acid, to not have oxidation products in the way, then heat?
There is also an equation for MnF3 hydrolysis: 2MnF3 + H2O => MnO2 + MnF2 + 2HF. Since in concentrated solutions with loads of HF that didproportionation hydrolysis doesn't happen, I have reason to believe it's reversible, otherwise MnF3 would eventually get totally hydrolysed, HF or no HF. Colud this be an even holier grail?
There are othe methods of chemical F2 production, will post them later...