Many substances that are generally regarded as unreactive, such as powdered steel, glass fragments, and asbestos fibers, are readily consumed by cold fluorine gas.

The noble metals ruthenium, rhodium, palladium, platinum, and gold react least readily, requiring pure fluorine gas at 300–450 °C (575–850 °F).

[17] The lighter noble gases xenon and krypton can be made to react with fluorine under special conditions, while argon will undergo chemical transformations only with hydrogen fluoride.

[18] Nitrogen, with its very stable triple bonds, requires electric discharge and high temperatures to combine with fluorine directly.

[27] As a result of its small size and high negative charge density, the fluoride anion is the "hardest" base (i.e., of low polarizability).

In the latter case, it significantly increases the acidity of a molecule: the anion formed after giving the proton off becomes stable as a result.

The explanation for the behavior is complicated, having to do with various cluster-forming tendencies of HF, water, and fluoride ion, as well as thermodynamic issues.

Zirconium, hafnium, plus many of the actinides form tetrafluorides with an ionic structure that puts the metal cation in an 8-coordinate square antiprism.

In this compound, manganese forms –MnF6– octahedra which share bridging fluorines to make –Mn4F20– rings which are then further connected three dimensionally.

The niobium and tantalum pentafluorides, have the same tetrahedra in their structures, with the difference being the formation of the tetra- (rather than poly-) meric molecules.

In combination with alkali metals, pentavalent bismuth can form hexafluorobismuthate, [BiF6]−, upon reaction with a fluoride donor, either strong (such as NaF[65][66]) or not (such as XeF4[67]).

[78] The reactivity of such species varies greatly—sulfur hexafluoride is inert, while chlorine trifluoride is extremely reactive—but there are some trends based on periodic table locations.

[81] SiF4 is stable against heating or electric spark, but reacts with water (even moist air), metals, and alkalis, thus demonstrating weak acidic character.

[2][88] The chalcogens (oxygen's periodic table column) are somewhat similar: The tetrafluorides of S, Se, and Te hydrolyze and are Lewis acidic.

They increase in reactivity with atomic number: SF6 is extremely inert, SeF6 is less noble (for example, reacts with ammonia at 200 °C (400 °F)), and TeF6 easily hydrolyzes to give an oxoacid.

The well-characterized heavier halogens (chlorine, bromine, and iodine) all form mono-, tri-, and pentafluorides: XF, XF3, and XF5.

Krypton tetrafluoride was reported in 1963,[104] but was subsequently shown to be a mistaken identification; the compound seems to be very hard to synthesize now (although even the hexafluoride may exist).

For example, the synthesis of mercury tetrafluoride, the first compound to achieve an oxidation state above +2 for a group 12 element, breaking the filled 5d-shell, again showing the significance of the relativistic effects on the heavy elements, and fueling the debate over whether mercury, cadmium, and zinc are transition metals,[110] occurred at cryogenic temperatures and the compound decomposes at the temperatures of solid nitrogen.

[111] More unstable still, a rare cobalt(V) species, the CoF+4 cation, has only been observed in gas phase (with no interactions with other atoms, thus no shown stability in any chemical environment).

Despite being often found as a short-lasting intermediate in the oxidation of water by fluorine, HOF can be still be stably isolated as a solid and the only hypohalous acid known to be able to do so.

After the reaction, the molecular size is not changed significantly, as the elements have very similar van der Waals radii.

[114] Direct fluorination becomes even less important when it comes to organohalogens or unsaturated compounds reactions, or when a perfluorocarbon is desired (then HF-based electrolysis is typically used).

[117] In contrast, the fluoropolymers are formed by polymerizing free radicals; other techniques used for hydrocarbon polymers do not work in that way with fluorine.

A vast number of small molecules exist with varying amounts of fluorine substitution, as well as many polymers—research into particular areas is driven by the commercial value of applications.

Substituting other halogens in combination with fluorine gives rise to chlorofluorocarbons (CFCs) or bromofluorocarbons (BFCs) and the like (if some hydrogen is retained, HCFCs and the like).

[120] However, if a perfluorocarbon contains double or triple bonds (perfluoroalkenes or -alkynes), a very reactive towards ligand accepting result, even less stable than corresponding hydrocarbons.

Perfluorinated compounds, as opposed to perfluorocarbons, is the term used for molecules that would be perfluorocarbons—only carbon and fluorine atoms—except for having an extra functional group (even though another definition exists[125]).

[127] These compounds lower surface energy; for this reason, they, especially perfluorooctanesulfonic acid (PFOS, formerly the active component in brand "Scotchgard") have found industrial use as surfactants.

[118] The simplest fluoroplastic is polytetrafluoroethylene (PTFE, DuPont brand Teflon), which is a simple linear chain polymer with the repeating structural unit:–CF2–.

The main challenges in making fluorelastomers are cross-linking (reacting the unreactive polymers), as well as removing the HF formed during curing.