It seems apparent, when one performs the standard "flame tests" in a lab, for the Group II elements that Magnesium burns by far the hottest and brightest. Why is this true?
I wouldn't expect this at all: as we descend Group II, we find elements with lower and lower electronegativities and greater electron shielding of the bonding pairs to the nucleus of the metal, so Barium should be much more reactive in forming oxides, halides etc. (this prediction is observed to be correct with, for example, the famous Group II reactions with water)
So one would assume that Magnesium would burn the least brightly of all the Group II elements, but instead we find that it burns white hot while the others you can look at in a dark room without damaging your eyes.
Now, I don't have any definitive data suggesting that Magnesium DOES burn at a greater temperature to the other Gr. IIs (nor could I find any online anywhere) so if I'm correct in the above assertion, why? And if not, why does it appear like that's the case?