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I was reading about the octet rule and saw exceptions in it. Different theories I read further and one of those was Sugden's concept of singlet linkage which state that the octet rule is never violated.

My question is that why Sugden's concept is not taught us or is not as popular as covalent bond or dative bond? Is there any problem with this concept? Does it have any limitations?

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    $\begingroup$ Hint: singlet linkage is a more accurate description than octet expansion, but the octet rule is still violated by octet-deficient molecules such as BeH2. $\endgroup$ – DHMO Oct 9 '16 at 16:07
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What he said

He did not say that the octet rule is never violated.

On the contrary, he suggested that "the maximum number of electrons in the valency orbit can not [sic] exceed eight". (source)

This means, while he denies that the octet rule is violated by molecules like chlorine pentafluoride, he did not challenge whether molecules traditionally labeled as octet-deficient such as beryllium difluoride indeed violates the octet rule.

He argued for the existence of a two-center one-electron bond.

Why his theory was poorly accepted

Preconceptions?

The success of his theory

To quote Wikipedia:

In the 1940s and 1950s, Rundle and Pimentel popularized the idea of the three-center four-electron bond, which is essentially the same concept which Sugden attempted to advance decades earlier; the three-center four-electron bond can be alternatively viewed as consisting of two collinear two-center one-electron bonds, with the remaining two nonbonding electrons localized to the ligands.

Indeed, the three-center-four-electron explains the bonding in the triiodide anion well.

Also, another success is that this theory, unlike the expanded-octet theory, does not rule out the existence of hypervalent second-period element, such as carbon ("Frozen" Transition States: Pentavalent Carbon et al., J. C. Martin, Science, 05 Aug 1983, Vol. 221, Issue 4610, pp. 509-514, DOI: 10.1126/science.221.4610.509).

The problems with this theory

He was reluctant to accept that ionic characters do indeed play a larger in some molecules traditionally labelled as hypervalent, such as the sulfate anion.

The failures of this theory

In the sulfate anion, there is no two-center one-electron bond (Wikipedia):

For sulfuric acid, computational analysis (with natural bond orbitals) confirms a clear positive charge on sulfur (theoretically +2.45) and a low 3d occupancy.

Also, some of the traditionally-labeled hypervalent molecules involve very electronegative elements such as fluorine for sulfur hexafluoride, oxygen for sulfate, and chlorine for phosphorus pentachloride. A better explanation for these molecules would need to invoke ionic characters instead of the covalent singlet linkage.

For sulfur hexafluoride (source):

The sulfur d orbitals are found to contribute very strongly to the binding energy ($250$ kcal/mol) and to have a total occupancy of around $0.25\mathrm{e}$, the valence electron distribution on sulfur being about $32\%$ in $3\mathrm{s}$, $59\%$ in $3\mathrm{p}$, $8\%$ in $3\mathrm{d}$, and $1\%$ in $4\mathrm{p}$. The occupancy of the two sulfur $3\mathrm{d}_{\sigma}$ orbitals ($0.16\mathrm{e}$), however, is only one-sixth of what would be required for $\mathrm{s}\mathrm{p}^3\mathrm{d}^2$ hybridization, and the energetic contribution of these orbitals is only two to three times larger than that of the sulfur $3\mathrm{d}_{\pi}$ orbitals. The sulfur $\mathrm{d}$ orbitals are important because they allow strong back transfer from the negatively charged fluorine ligands to the strongly positively charged ($+2.9\mathrm{e}$) sulfur, in turn allowing significant contraction of the $\ce{S-F}$ bonds and greatly increased molecular stability.

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