According to Wikipedia, tellurium tetraiodide decomposes on heating:

$$\ce{TeI4 ->[\Delta] TeI2 + I2}$$

However, the reaction is somewhat wrong as according to this 2008 paper:

Solid tellurium tetraiodide decomposes to form $\ce{I2(g)}$ and $\ce{Te(s)}$. A very small contribution of $\pu{3.3 ± 2.1 mol\%}$ of gaseous $\ce{TeI2}$ was also determined by both GED and MS. The "metallic" $\ce{Te}$ accumulated in the solid phase vaporizes at above ca. $\pu{670 K}$ as the predominately $\ce{Te2}$ molecular species

enter image description here

In the other Wikipedia article, it says that tellurium diiodide has never been isolated in bulk although organic complexes are well known and characterized.

Why has tellurium diiodide never been isolated in bulk? Is it unstable/prone to disproportionation to monoiodide/subiodide? We never got to know its structure nor its physical properties like b.p, m.p, density, color etc. other than the fact that the species is gaseous?

Question: Has tellurium(II) iodide been properly characterized ever?

  • $\begingroup$ Just curious, what is interesting about this specific compound? Does your university subscribe to SciFinder and/ or Reaxys? SciFinder shows that in the last 100 years there are only ~35 papers /patents/reports that even mention it. However, my point is that there are perhaps millions of other structures which do not exist...so back to the original question, what is interesting about it? $\endgroup$
    – AChem
    May 14, 2022 at 1:21

1 Answer 1


It isn't just tellurium(II) iodide. Lower chalcogen halides tend to disfavor the chalcogen(II) compound in favor of subhalides where chalcogen atoms are bonded to each other and have oxidation state no greater than +1.

Let us look at the sulfur-chlorine system, where both the "normal" halide $\ce{SCl2}$ and the subhalide $\ce{S2Cl2}$ are known. However, "sulfur dichloride loses chlorine slowly at room temperature and reverts to disulfur dichloride."[1].

The preference for disulfur dichloride in this system may be traced to pi backbonding. The structure of disulfur dichloride is given here from Ref. 2:

enter image description here

With this structure, a pi-symmetry $p$ orbital on sulfur, which is a good pi donor, overlaps with the conjugated sulfur-chlorine antibonding orbital leading the an electron transfer indicated by the blue arrow, weakening the sulfur-chlorine sigma bond but creating a pi bonding interaction between the sulfur atoms:

enter image description here

In effect, what is formally the sulfur-chlorine bond is delocalized into the sulfur-sulfur pi orbitals. A similar interaction involving the second chlorine atom occurs on the orthogonal plane (note the near-right dihedral angle).

The impact of this interaction is seen by comparing bond lengths. The sulfur-chlorine bond length shown above is $206$ pm, longer (meaning weaker) than the $201$ pm given for monosulfur dichloride in 1. The sulfur-sulfur bond length, by contrast, at $195$ pm is shorter (thus stronger) than $205$ pm on molecular octasulfur [3]. While a similar interaction is possible with $\ce{SCl2}$, it would require chlorine instead of sulfur to donate the electron pair and take the positive formal charge, rendering the backbonding less favorable.

In the case of lower tellurium iodides, there are two preferred subhalide stoichiometries, $\ce{TeI}$ (two phases) and $\ce{Te2I}$ [4]. These have polymeric structures in which the tellurium forms chains, allowing for further bond delocalization through the pi backbonding.


  1. https://en.wikipedia.org/wiki/Sulfur_dichloride, retrieved 08 May 2022.

  2. https://en.wikipedia.org/wiki/Disulfur_dichloride, Retrieved 08 May 2022.

  3. https://en.wikipedia.org/wiki/Octasulfur, retrieved 08 May 2022.

  4. https://en.wikipedia.org/wiki/Tellurium_iodide, retrieved 08 May 2022.


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