Group I/II Metal oxides always basic in aqueous solution -- why?

[adf89812]

TL;DR Summary

Make argument rigorous

Transform my handy-wavy explanation of why group I/II metals oxides in aqueous solution are basic into a rigorous argument​

I heard an explanation about something being a better proton acceptor or lone pair donor but that doesn't make sense. I couldn't explain in in terms of acid-base theory.

The hand-waving way I saw it was that group I/II metal oxides are less electronegative than non-metals, so in water, they'll donate their electron to the hydrogen, the hydrogen will break away from the oxygen because hydrogen hates oxygen hogging its electrons, and because hydrogen electronegative enough.

With non-metals, my hand-waving is that metal oxides are more electronegative, when they bond with water, they'll just form one bigger molecule because they suck on other's electrons without letting go and form one big acid molecule where the least electronegative thing in there is a hydrogen, which falls of into a proton, and it may or may not be polyprotic.

 

[Borek]

group I/II metal oxides are less electronegative than non-metals

Electronegativity is a property of an element, not of a compound.

 

[adf89812]

Electronegativity is a property of an element, not of a compound.

electronegative doesn't disappear when you form a chemical bond. Oxygen doesn't become less electronegative AFAIK when it bonds.

 

[adf89812]

electronegative doesn't disappear when you form a chemical bond. Oxygen doesn't become less electronegative AFAIK when it bonds.

Maybe there's something about net electronegativity effect in molecules?