Fluorine is the most electronegative element. That means it will terrorise other elements for an electron to achieve the electron configuration of neon.
As a matter of fact it cannot even cohabit with itself. The F−F bond in the fluorine molecule is the weakest of the halogen gases and so it is the most reactive of the gases. In chemistry term we say that there is serious repulsion between the nonbonding electrons because of the short bond length of the F2 molecule.
So fluorine is the most reactive element in chemistry. It either forms an anion F‾ or a covalent bond.
With the metallic elements it form a salt. The radius of F‾, 1.36�, is almost identical with that of O‾� (1.40�). Fluorine is found in nature as insoluble fluorides; notably fluorspar, CaF2, cryolite, Na3AlF6, and fluorapatite (a complex of Ca with phosphate, fluoride and chloride). The amount of fluorine in the earth's crust is about the same as chlorine, about 0.06%. The only stable isotope is 19F.
With less electropositive elements it forms covalent fluorides. Oxygen is an exception as it is also very electronegative. The other important factor is the small size of the fluorine atom. This allows it to get very close to the other atom to form a exceptionally strong covalent bond. Just remember Teflon used in non-stick frying pans and cooking pots. The bond with carbon is so strong that it can withstand high heat without breaking up.
HYDROGEN FLUORIDE
Hydrogen fluoride is made by the reaction of sulphuric acid on CaF2 and is the principal starting reagent for fluorine and fluoride compounds. The fluorides of the metallic elements − Group IA, Group IIA and Group IIIB - are prepared by reacting the metallic oxides or carbonates with HF. The other compounds are prepared by heating the respectively chloride with CaF2 or direct fluorination of the element.
HF can attack glass, so it is best handled using metal or plastic equipment.
The bond between hydrogen and fluorine is highly polar; even ionic. After water, liquid HF (at low temperature) is one of the most important solvents. It dissolves both inorganic and organic compounds. In liquid form it is actually a mixture of fluoride complexes; [H2F]+, F‾ and [HF2]‾ ions.
When dissolve in water it is a weak acid
The fluoride ion, as a conjugate base, can react with HF;
FLUORIDES
The fluorides of the alkali metals are soluble in water, however lithium fluoride is only slightly soluble in water and insoluble in alcohol. This is most likely due to the intense electrostatic attraction between each other. The charge per unit surface area is very high owing to the small size of the ions.
In Group IIA, beryllium fluoride shows partial covalent bonding and is readily soluble in water. The other congeners form ionic compounds with fluorine. Magnesium fluoride is sparingly soluble in water, but the solubility increases with increasing size of the element starting from magnesium fluoride.
In Group IIIB the size effect of fluorine (getting closer to the other atom) is becoming more significant. The bond length between the boron and fluoride atom is shorter than expected for a sigma bond. A possible explanation is the addition bonding between the filled p orbital of the fluorine with the empty p orbital of the boron atom.
Boron fluoride is the strongest Lewis acid known.
The fluorides of Al, Ga, and In are ionic and extremely stable. They can be heated to high temperature (about 1000�C). This is because the small size of the fluoride allows the metal to attain a maximum coordination of six and so it tends to have large stable crystal structures.
The elements also form various complexes with fluorine. Aluminum has six fluoride complexes, from [AlF]‾� to [AlF6]‾�.
In Group IVB the formation of F‾ complexes is even more dominanting. Silicon and germanium tetrafluorides will give the hydrous oxide and the hexafluorides, [SiF6]‾ and [GeF6]‾, with water.
The other silicon halides will dissolve in water to give the silicic acid. Tin do form a tin difluoride, SnF2. This difluoride gives weak complexes with compounds have oxygen of nitrogen atom by interacting with the lone electron pair.
In Group VB the elements form two types of binary fluorides − MF3 and MF5. PF3 is not easily hydrolysed by water and do not form F‾ complexes. The other fluorides are rapidly hydrolysed by water and form complex fluoride ions like [SbF2]+, [SbF4]‾, and [SbF5]‾�. PF5 is a very strong Lewis acid. It can form a [PF6]‾ complex with HF. Many hexafluorophosphate salts are known. AsF5 and SbF5 behave like PF5. SbF5 is a very powerful fluorinating agent.
In Group VI, the electronegativity of oxygen is sufficient high and so it is not in favour of sharing electrons with fluorine to form a bond. It is the only element that does not form a binary fluoride. With the other congeners, fluorine forms two types of binary fluoride - the hexafluorides and the tetrafluorides.
Sulphur hexafluoride is very stable and very resistant to chemical attack. Selenium and tellurium hexafluorides are more reactive. Tellurium hexafluoride has weak Lewis acidity and is fully hydrolysed by water very slowly.
SF4 is an extremely reactive. It is hydrolysed by water to SO2 and HF instantly. Selenium and tellurium tetrafluorides are also reactive. All three are good fluorinating agents.
