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It is known that most dianions are unbound per se, i.e. that the corresponding monoanions have negative electron affinities, and the "dianions" we see in e.g. metal oxides can exist as such only due to the additional positive charges (on top of the "weaker" positive charges on the e.g. oxygen cores) on the countercations' stabilising effect on the dianion's electrons, rendering them bound. It is also known that most neutral molecules are electronically bound; i.e. that the corresponding monocations have positive electron affinities.

However, some monoanions do have positive electron affinities- for example, the dodecahydrido-closo-dodecaborate dianion is bound, rendering its (radical) monoanion's electron affinity being positive.

My question now follows- are there any monocations with negative electron affinities; i.e. whose neutral counterparts are electronically unbound?

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    $\begingroup$ I am not aware of any atom or molecule with negative ionization energy. Probably because their bound electrons have lower energy then free electrons. $\endgroup$
    – Poutnik
    Jul 20 at 6:20
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    $\begingroup$ Are you asking whether there are any electronically unstable atoms prone to spontaneously lose an electron and form stable cations? In the absence of a solvent? $\endgroup$
    – Buck Thorn
    Jul 20 at 6:30
  • $\begingroup$ @BuckThorn yes. $\endgroup$ Jul 20 at 7:24
  • $\begingroup$ Atoms OR neutral molecules. $\endgroup$ Jul 20 at 7:25
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    $\begingroup$ Better question would be if it's possible at all, because it actually may be, even if the whole idea is way more outlandish than it may seem. $\endgroup$
    – Mithoron
    Jul 20 at 17:30

2 Answers 2

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A monocation with an endothermic (negative) electron affinity is the same as a neutral species with an exothermic (negative) ionization energy. Reframing it this way may turn up more information.

The neutral substance with the lowest ionization energy I know is tetrakis(hexahydropyrimidinopyrimidine) ditungsten(II), also called ditungsten tetra(hpp). With a positive ionization energy of +3.51 eV, it is still very, very far from the goal. It does not seem possible that a reasonable chemical system would manage to bring the ionization energy all the way down below 0 eV. The problem is that ionization energy/electron affinity are defined and measured in a vacuum, but opposite net charges in a vacuum are completely unscreened and attract far too strongly.

But what if you do screen the charges? In a much weaker sense, you might consider that electrides kind of fit the requirement - you can think of these as neutral compounds which, in a condensed phase (typically the solid state), spontaneously eject an electron as if it were a separate anion. The condensed phase surrounds the charges and does the job of screening them enough so that the electrons and the cations can exist as "separate" entities.

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  • $\begingroup$ Meh, I never liked such view on electrides. $\endgroup$
    – Mithoron
    Jul 20 at 20:11
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You are talking about superalkalis, which are defined as groups of atoms that have a lower ionization energy than any alkali metal atom -- that limit being about 3.9 eV for caesium. ("Pseudo-alkali" is used for iniization energies below the 5.4 eV value for lithium.)

As noted by Nicolau Saker Neto, no superalkali is known with ionization energy less than 3.5 eV, so we do not have anything that spontaneously gives off an electron to form a gas-phase superalkali electride. However, lattice-energy stabilization effects do stabilize some condensed electrides, such as the cryptand-alkali metal electrides given in Ref. [1].

Some superalkalis have a lower ionization energy than the electron affinity of atomic chlorine (which actually beats that of fluorine). Trilithium oxide, $\ce{Li3O}$, is one example; the species identified by Nicolau is another. For such species $\ce{M}$, the combination $\ce{M^+, Cl^-}$ would be more stable, without ion-pairing or condensation, than the neutral $\ce{M}$ group plus a neutral chlorine atom. This comparison does not, however, account for the bonding of chlorine in its normal state into diatomic molecules.

References

  1. James L. Dye (19900). "Electrides: Ionic Salts with Electrons as the Anions". Science 247, 4993, 663-668.
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