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The large drop from carbon to silicon is the main trend; the smaller reversals at Ge and Pb reflect changes in shielding and valence-electron binding. Allen electronegativity is based on the mean binding energy of valence electrons in free ground-state atoms, rather than an arbitrary charge-transfer count. Moving to larger principal shells generally makes the valence electrons more diffuse and more shielded, reducing their attraction to the nucleus. carbon's compact second-shell valence orbitals therefore give particularly high electronegativity.

The intervening filled 3d shell shields imperfectly, increasing effective nuclear attraction in germanium relative to a simple smooth extrapolation from silicon. This helps explain Ge being slightly above Si rather than below it. Increased size then lowers Sn relative to Ge. The filled inner d and f shells and relativistic stabilization of lead's 6s electrons modify the next step; contraction and stronger binding of the 6s contribution help Pb lie slightly above Sn. The small reversals should not obscure the much larger C-to-Si change. The spectroscopic definition includes both occupied s and p valence contributions, so the relativistic s effect can influence the Allen average even though not every electron participates equally in every bond.^[ref]

The central carbon is approximately $sp$ hybridized, with two opposite sigma bonds and no carbon lone pair; the electron-rich donor ends are sulfur and nitrogen. The 16 valence electrons allow the principal resonance structuresEach has octets and a linear S–C–N skeleton. Carbon's two perpendicular p orbitals support the pi interactions. The first contributor emphasizes the C≡N bond and sulfur-centred negative charge; the second distributes pi bonding over S=C=N and puts negative charge on N. The real anion has partial multiple-bond character and delocalized charge rather than alternating between these drawings.

Nitrogen's greater Allen electronegativity favours electron accumulation there, while sulfur supports a diffuse, polarizable electron-rich region. Thus both terminal atoms are important electron-rich sites relative to central carbon. “Greatest electron density” should not be interpreted as a literal comparison of all core electrons on different elements; the chemically relevant comparison is valence charge, donor density or electrostatic potential. The exact partition of negative charge between S and N is not uniquely determined by electronegativities or Lewis structures alone.

Allen electronegativity is proportional to the average binding energy of the free ground-state atom's valence electrons. For an s,p main-group atom, use positive, spectroscopically inferred one-electron binding energies ${\epsilon }_s,{\epsilon }_p$ and their occupancies $n_s,n_p$:Here $C>0$ fixes the conventional scale; the one-electron energies are obtained from appropriate atomic ionization/term energies, not merely from the first ionization energy. For group 15 the weighted average is $(2{\epsilon }_s+3{\epsilon }_p)/5$.^[ref]

The approximate connection to effective nuclear charge is ${\epsilon }_{n\ell }\sim R_{\infty }[Z_{\mathrm{eff}}(n\ell )/n_{\mathrm{eff}}{]}^2$. At a fixed shell, stronger effective nuclear attraction generally raises the valence binding energies and therefore Allen electronegativity. Down a group, increasing shell size and changes in penetration offset that attraction. There is no universal relation ${\chi }_{\mathrm{Allen}}\propto Z_{\mathrm{eff}}$ across different shells, nor can one substitute a core-electron $Z_{\mathrm{eff}}$ for the valence average.

Sulfur substituting for phosphorus acts as a donor impurity, giving n-type semiconductor behaviour and usually increasing electrical conductivity. A P site normally supplies five valence electrons; S supplies six. After forming the surrounding bonds, the extra electron can occupy a donor level below the conduction band and be thermally ionized into that band. The donor impurity then carries positive charge.

This assignment is supported experimentally: infrared transmission measurements on S-doped gallium phosphide resolved bound-donor transitions and determined a sulfur donor optical ionization energy of about $0.104\mathrm{eV}$, far smaller than the host band gap.^[ref]

The effect depends on substitution on P sites, donor ionization, compensation by acceptors and charge carrier mobility. A dilute active donor population increases the electron term in $\sigma =e(n{\mu }_n+p{\mu }_p)$; not every dopant must donate a mobile electron at room temperature. Sulfur's larger Allen electronegativity does not overturn this electron-counting argument or make a P-site sulfur atom an acceptor.