According to J.D Lee, compounds with fraction bond number are unstable. I calculated that the bond order of $\ce{Li2^2+}$ is 0.5 while that of $\ce{Li2}$ is 1. Hence, $\ce{Li2^2+}$ must be less stable than $\ce{Li2}$ due to half bond character.

But, in reality, $\ce{Li2}$ is more stable than $\ce{Li2^2+}$. Why is it so?

  • $\begingroup$ Your question makes no sense: "Hence Li2+ must be unstable than Li2 but then why Li2 is more stable than Li2+" says the same thing twice. What are you trying to ask? $\endgroup$
    – Ian Bush
    Jun 10 '18 at 6:31
  • $\begingroup$ i m trying to ask why li2+ is more stable than li2? $\endgroup$
    – Hercules
    Jun 10 '18 at 6:34
  • $\begingroup$ I suggest you edit your question to say that then. Further stable with respect to what? The dissociation products of the two molecules are different. Do you mean bond strength? I also suggest you think more carefully about the tags you use for your question, organic-chemistry, periodic-table and electronegativity are all totally irrelevant. $\endgroup$
    – Ian Bush
    Jun 10 '18 at 9:23
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    $\begingroup$ Rafael, please don't rollback edits that attempt to improve your post. If you have some reasons to completely revert the changes, please also let the particular editor know that you're reverting their edits via the @ notification. Thanks. $\endgroup$ Jun 10 '18 at 14:17
  • $\begingroup$ @GaurangTandon He's asking about bond order 1/2 so definitely not +2 charge. $\endgroup$
    – Mithoron
    Jun 10 '18 at 19:11

The comment asks "Why is Li2+ more stable than Li2". Li2 has a relatively low bond energy (gas phase) of 27 kcal/mol (Cotton and Wilkinson, Inorganic Chemistry). The energy cost to make one mole of Li gas is 37 kcal (CRC Handbook). The ionization energy of Li is 5.39 eV = 124 kcal/mol.

So, breaking a Li2 gas molecule into 2 Li (gas) costs 27 kcal/mol.

Breaking a Li2+ gas molecule into Li + Li+ involves not the separation of two uncharged atom, but the separation of a very small Li+ ion from a Li atom which provides some charge accomodation. The heat of formation of Li+ in water is 66 kcal/mol (presumably from aquation with 4 waters); the hydration bonding is worth about (124 + 37 + 66)/4 = 54 kcal/mole each. So I would estimate the bond between Li+ and Li to be perhaps 40 kcal/mol, at least a significant fraction of 54 kcal/mol, and probably more than 27 kcal/mol.

The important point is that the unusually small size of the Li+ ion makes it able to polarize a neutral atom to form a strong bond (and come in closer!), whereas the neutral atoms have only uncharged molecular orbitals to spread their electrons over. It is highly unlikely that larger atoms (Na, K) would show similar stability for an ion-molecule compared to a neutral molecule. Magnesium might be a similar exception, however, since both Li and Mg have small radii; it might be interesting to compare stabilities of Mg2 and Mg2+.


Here you're saying that Li2+ is more unstable than Li2

Hence Li2+ must be unstable than Li2

And here you're confirming that Li2 is more stable than Li2+, corroborating with your previous statement

but then why Li2 is more stable than Li2+

But in your question you're saying the opposite, that Li2+ is more stable than Li2. So i'm confused with what you're trying to ask.

In my knowledge, based on inorganic text books, you can explain the stability of a molecule using molecular orbital theory (MOT). But MOT is a bond theory and Li2+ has no bond unless you're talking about (Li-Li)2+. If this is the case, Li2 is more stable than (Li-Li)2+ -(Li-Li)2+ is not even going to exist according with MOT - since the bond order for Li2 is 1 and (Li-Li)2+ is 0 as shown in the diagram below.

enter image description here

If you're trying to compare Li2+ and Li2 I don't know if it's possible, i've never seen something like that.

  • $\begingroup$ He's probably asking about $\ce{Li2+}$. If you don't understand question then why to answer? (Rhetoric question) $\endgroup$
    – Mithoron
    Jun 10 '18 at 19:09
  • $\begingroup$ Someone commented a question about my answer and deleted after a while. I didn't see that the person deleted it. $\endgroup$ Jun 10 '18 at 19:18

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