Imagine an inorganic complex with a bunch of terminal ligands. All else the same, would $\ce{F-}$ or $\ce{Br-}$ ligands be more electron-donating and why?

Of course, in terms of $\mathrm{p}K_{\mathrm{a}}$, $\ce{F-}$ is more basic than $\ce{Br-}$, predominantly due to the smaller size of the $\ce{F-}$ anion. I have also read in multiple places that the $\mathrm{p}K_{\mathrm{a}}$ trends can be used to describe Lewis acidity/basicity trends, although it's not immediately clear to me why this is true. If one takes the prior statement at face-value, then it appears that that $\ce{F-}$ is more Lewis basic and, by the very definition of a Lewis base, is more electron-donating than $\ce{Br-}$. However, why would $\ce{F-}$ be more electron-donating if its electronegativity is higher than $\ce{Br-}$? These two points seem paradoxical.

For full context, I am trying to understand why $\ce{F-}$ would be better at stabilizing higher oxidation states of metals in inorganic complexes than $\ce{Br-}$, which based on the literature appears to be related to the electron-donating strength of the ligands (with more electron-donating ligands increasing the stability of higher oxidation states).


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