Relationship between Lewis Theory and Redox theory? First of all, thank you ahead of time for taking the time to read my question and answer! 

Let us take the reaction $\ce{SnI4 + I^{-} -&gt; SnI2 + I3^{-}}$ as an example.

Here we see one obvious redox reaction, $\ce{Sn(IV)}$ gets reduced to $\ce{Sn(II)}$.
Less obviously, the iodide gets oxidized to $\ce{I(I)}$ in the triiodide anion. The other 4 iodines maintain their negative 1 oxidation state. 

Is there also a lewis theory acid/base reaction here as well?
The tin acts as a lewis acid and accepts an electron from each of the two of the iodines (lewis bases)? Those iodines, now electron deficient (act as lewis acids), bond with the iodide (acting as a lewis base) to create $\ce{I3^{-}}$ product.

In this case, Tin accepts one electron from two atoms rather than a "pair of electrons" from one atom like you see in traditional examples of lewis acids.
So that would be my first question, can Tin act as a "lewis acid" even though it is accepting one electron from two different sources rather than a pair from one source? Or in other words, can the iodines be considered "lewis bases" even though they only donate one electron?

Or is the whole paragraph above a moot point because there is no acid/base reaction here because the Tin is simply regaining "ownership" of the electrons it let iodine "borrow" to create a covalent bond and thus its not actually getting any "new" electrons. 

My second question is: Can we make a general statement about the relationship between lewis acids and their roles in a redox reaction? For example, can I say that ALL lewis acids are oxidants and ALL lewis bases are reductants? Or are these concepts mutually exclusive? In other words, sometimes lewis acids can act oxidants, sometimes as reductants, and sometimes as neither, like in the example below?:

$\ce{Al(III)Cl3 + Cl- -&gt; Al(III)Cl4-}$ (no redox here)
Under the Usanovich definition of acids and bases, any substance which loses electrons is a base while any substance which gains electrons is an acid. Here, any redox reaction is also an acid-base reaction, whereas Lewis theory has nothing to state in these cases. 

Regarding your second question, blanket statements are dangerous. Depending on exactly what you mean, I can provide a counter-example. Here is a strong oxidant ($\ce{HNO3}$) reacting as a Lewis base:

$$\ce{HNO3(l) + H2SO4(l) -&gt; H2NO3^{+} + HSO4^{-}}$$