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?

  • $\begingroup$ Here's a start: chemguide.co.uk/inorganic/group2/reacto2.html $\endgroup$ – Martin Jan 26 at 21:59
  • 1
    $\begingroup$ "But how reactive a metal seems to be depends on how fast the reaction happens - not the overall amount of heat evolved. The speed is controlled by factors like the presence of surface coatings on the metal and the size of the activation energy." $\endgroup$ – Martin Jan 26 at 22:02

Your Answer

By clicking “Post Your Answer”, you agree to our terms of service, privacy policy and cookie policy

Browse other questions tagged or ask your own question.