Does the charge on an atom impact the polarization of covalent bonds?

So far, I have generally seen polarization power only discussed in terms of ionic bonds: small, highly charged cations have more polarization power, and large anions are more polarizable. The more polarization power and polarizability, the greater the covalent character of an ionic bond.


Generally, one can estimate the polarity of a covalent bond by looking at the differences in the electronegativities of the two atoms involved in the bond. However, I am wondering whether one can apply the ideas of polarizing power to atoms in covalent bonds, specifically the idea that net charge impacts the polarity of the bond. Does a higher positive net charge on one atom cause electrons in covalent bonds to be pulled towards that atom? This is what I would conclude from a polarization power argument, because an atom with a positive net charge has more polarization power. Similarly, would an atom with a negative charge pull less electrons? These ideas also seem to coincide with a Coulomb's Law argument that a more positive charge would attract electrons better, and a more negative charge would not attract them well.


One piece of evidence supporting this idea is the inductive effect; polarization propagates down a chain of sigma bonds because partially positive carbons are able to polarize carbon-carbon bonds. This polarization occurs in spite of the fact that the difference in electronegativity between the two carbon atoms is 0. This seems to imply that one has to consider both electronegativity and the charge on the atoms when determining the polarity of bonds. [![enter image description here](https://i.stack.imgur.com/GESmV.png)](https://i.stack.imgur.com/GESmV.png)


There are two particular cases that I am curious about: metals in coordination complexes and formal charges. Consider two complexes: one with a Mn (II) center and one with a Mn (VII) center. Will the Mn (VII) center withdraw the electrons of the M-L dative bonds better than an Mn (II) center (in turn inducing an electron withdrawal throughout the rest of the ligands)? For example, would an NH2 ligand be less basic when bonded with an Mn (VII) metal center than when bonded to an Mn (II) metal center, because the Mn (VII) center is able to withdraw electrons better (even though an Mn-N bond, with the same electronegativity difference between Mn and N, exists in both cases)?


As for formal charges, would an atom with a positive formal charge (such as the nitrogen in NH4+) withdraw electrons better than one with a negative formal charge (such as the nitrogen in NH (2-))? I know that formal charge does not always represent the true charge on an atom, but I am thinking that it could be used to compare the relative amounts of charge on related species - for example, I know most electron density is concentrated on the nitrogen in both NH3 and NH4+, but I assume that there would be a bit less electron density on the nitrogen in the NH4+ because the lone pair electrons of NH3 are now shared with a hydrogen.


I am also aware that there are other factors to consider when looking at the electron distributions of complexes and formal charge, such as pi donors/acceptor ligands and favored resonance states. However, I am wondering about the effects of charge on the polarization of bonds. Thank you for your help.




Everything is correct here, including the explicit examples- the ammonium group is a very strongly inductively electron withdrawing group in electrophilic aromatic subtitution, far negating any hyperconjugative mesomeric electron donation effects, and "permanganamide" is likely to be strongly acidic, with the Mn-O and Mn-N bonds being highly covalent.


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