Can fluorine form triple bonds? A plain-language look

Fluorine is the most electronegative element in the periodic table. In plain words, it pulls electrons toward itself harder than any other atom. That is why textbooks treat it as a pure electron acceptor. It takes electrons and does not give them away.

A bond is two atoms sharing electrons. A single bond shares one pair, a double bond two pairs and a triple bond three. Nitrogen and carbon form triple bonds all the time. Fluorine is usually drawn with exactly one bond.

Some chemists know that fluorine can bond both ways. It can act as an acceptor and as a donor of electrons. The paper points to the small molecule NF, where the bond between nitrogen and fluorine already has partial double-bond character.

If fluorine can go to partial double bonds, the next question is obvious. Can it go further? Other elements in the same row of the periodic table form triple bonds. The preprint asks whether fluorine can too.

The study is theoretical. It uses quantum chemistry calculations, which predict how electrons behave in a molecule, and no lab experiments.

The main test case is MnF, a pair of one manganese atom and one fluorine atom. Manganese is a transition metal with many electrons that can take part in bonding. The work asks whether fluorine and manganese can form a triple bond in this molecule.

It then widens the view. It looks at other single fluorides that have the same number of electrons as MnF, at MnF in higher oxidation states (where manganese has given away more electrons), and at heavier transition metals paired with halogens, the fluorine family. For these it checks how likely the molecules are to form singlet states, an electron arrangement that can point to multiple bonds.

To push fluorine further toward giving electrons, the study also looks at more highly charged, electron-rich metals. It names Ni(III)F and Pd(III)F, nickel and palladium fluorides with a charge of +3 on the metal, as likely candidates for a fluorine triple bond and calculates both.

The paper draws two conclusions. First, the Wigner-Witmer rules, a simple way to predict the electronic states of a two-atom molecule from its atoms, may reach their limit for MnF, molecules with the same electron count and heavier relatives.

The clearest case is nickel. Ni(III)F binds fluorine much more strongly than manganese does, which the paper explains by the positive charge letting nickel use fluorine's donor side. Its calculated bonding orbitals show triple-bond character from fluorine to Ni(III). Pd(III)F should sit in between, but there the ECP shortcut gave unreliable results.

Second, it shows that effective core potentials can give misleading results for heavier transition metals bonded to halogens. Effective core potentials (ECPs) are a common shortcut. They replace the inner electrons of a heavy atom with a simpler stand-in to save computing time. The paper argues that this shortcut can go wrong in exactly these molecules.

For people outside chemistry, the takeaway is simple. A fact everyone learns, that fluorine only ever forms one bond, may not hold in every molecule. And two standard tools chemists use to check such questions can give the wrong answer here.

Electronegativity: how strongly an atom pulls shared electrons toward itself. Fluorine has the highest value of all elements.

Donor and acceptor: an electron donor shares electrons into a bond, an acceptor takes them in. The preprint asks how far fluorine can act as a donor.

Transition metal: elements like manganese, nickel and palladium in the middle block of the periodic table. Their many outer electrons allow complex bonding.

Oxidation state: a bookkeeping number for how many electrons an atom has given away or taken in a compound.

Singlet state: an electron arrangement in which all electron spins are paired. The preprint uses the tendency to form singlets as a hint for multiple bonds.

Halogens: the fluorine family, which also includes chlorine, bromine and iodine.

This is a computational preprint and has not been peer reviewed. It does not report a lab measurement of a fluorine triple bond. To test the idea in the lab, it suggests trying Ni(III)F or Pd(III)F, for example stabilised with weakly coordinating anions and studied by spectroscopy. Its conclusions are about what calculations suggest and where common methods break, not a final answer to the title question.

The same question about molecules near the limits of standard methods comes back in Julian M. Kleber's later work on the CARATE molecule model, which cites this preprint.